UNIVERSITY  OF  CALIFORNIA 
DEPARTMENT  OF  CIVIL  ENGINEER^ 
PERKELEY.  CALIFORNIA 


Engineering 
Library 


UN.VERS.TY  OF  CAUFORN.A 


OF 


'ERKELEY.  CALIFORNIA 


* 
4* 


AN  INTRODUCTORY   COURSE  IN 

QUANTITATIVE   CHEMICAL   ANALYSIS 


THE  MACMILLAN  COMPANY 

NEW  YORK  •    BOSTON  •   CHICAGO  •   DALLAS 
ATLANTA  •   SAN  FRANCISCO 

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THE  MACMILLAN  CO.  OF  CANADA,  LTD. 

TORONTO 


AN  INTRODUCTORY  COURSE  IN 

I"    QUANTITATIVE 
CHEMICAL   ANALYSIS 


WITH 


EXPLANATORY  NOTES,  STOICHIOMETRICAL 
PROBLEMS  AND  QUESTIONS 


BY 


GEORGE  McPHAIL  SMITH 

n 

Associate  Professor  of  Chemistry  in  the       t 
University  of  Illinois 


'        o  .   .   ,  >     „       •>,       »        .    .       , 

5  '»   '       JO»J     >'*"*»»" 


Ntfo  gorfc 

THE  MACMILLAN  COMPANY 
1919 

All  rights  reserved 


S5" 


Engineering 
Library 


COPYRIGHT,  1919, 
BY  THE  MACMILLAN  COMPANY. 


Set  up  and  electrotyped.    Published  July,  1919. 


J.  S.  Gushing  Co.  —  Berwick  &  Smith  Co. 
Norwood,  Mass.,  U.S.A. 


PREFACE 

THIS  introductory  course  in  Quantitative  Analysis  is  designed 
for  use  with  classes  consisting  of  students  who  have  completed 
courses  in  Elementary  Chemistry  and  Qualitative  Analysis,  and 
who  are  beginning  work  in  Quantitative  Analysis.  On  the 
laboratory  side,  its  primary  intent  is  to  provide  the  student  with 
directions  sufficiently  detailed  to  offer  little  opportunity  for 
going  astray,  and  thus  to  enable  him  to  work  successfully  without 
an  undue  amount  of  personal  supervision.  The  instructor  is 
thereby  placed  in  a  position,  in  the  laboratory  as  well  as  in  the 
classroom,  to  exert  his  personal  influence  more  especially  towards 
the  development  of  theoretical  knowledge  and  independent 
thought  on  the  part  of  the  students. 

The  use  of  the  book  in  the  laboratory  should  of  course  be 
supplemented  by  regular  classroom  instruction ;  and,  with  this 
in  mind,  it  has  seemed  desirable  to  include  the  stoichiometrical 
problems  of  Part  IV,  and  the  questions  of  Part  V.  The  problems 
of  Part  IV  are  such  as  are  constantly  met  with  in  analytical 
work,  and  their  conscientious  study  will  give  the  student  an  in- 
sight into  the  principles  of  a  wide  variety  of  processes;  the 
answers  to  the  problems  have  been  intentionally  omitted.  It  is 
the  writer's  practice  to  require,  as  a  written  exercise  to  be  handed 
in  at  the  beginning  of  a  recitation,  the  solution  of  a  definite 
number  of  problems  each  week  throughout  the  course ;  these  are 
graded,  and  are  returned  at  the  end  of  the  following  recitation. 
The  questions  of  Part  V  are  for  the  most  part  answered  in  the 
notes  or  elsewhere  in  the  book;  but  it  has  been  the  writer's 
experience  that  the  beginner  reacts  more  favorably  to  concrete 
questions  assigned  in  advance  for  study,  than  to  the  same  ques- 
tions when  put  to  him  for  the  first  time  just  after  he  is  supposed 


vi  PREFACE 

to  have  mastered  the  principles  and  details  of  a  specific  analytical 
process. 

The  general  directions  and  discussions  of  Part  I  are  intended 
to  emphasize  those  matters,  both  of  theory  and  practice,  which 
should  receive  especial  attention  from  the  worker  in  analytical 
chemistry.  It  is  of  course  realized  that  a  mere  reading  of 
Part  I  will  not  go  far  towards  familiarizing  the  student  with  its 
contents ;  but  it  is  sought  to  accomplish  this  end  by  referring 
later  on  in  the  text  to  special  subjects  as  occasion  presents. 

The  analyses  selected  for  practice,  included  in  Parts  II  and 
III,  are  those  which  are  comprised  in  the  elementary  courses  of 
quantitative  analysis  at  the  University  of  Illinois.  They  have 
been  chosen  as  being  satisfactory  types  of  gravimetric  and 
volumetric  analysis,  and,  after  several  years'  experience,  they 
are  considered  to  afford  to  all  classes  of  students  a  suitable 
foundation  for  more  advanced  work.  It  is  believed  that  they 
furnish  also  a  good  insight  into  the  methods  of  quantitative 
analysis,  and  hence  are  adapted  to  the  needs  of  students  who 
will  not  extend  their  study  beyond  the  period  of  an  introductory 
course. 

In  addition  to  the  help  derived  from  other  books  and  from 
journal  articles,  the  writer  wishes  to  acknowledge  his  especial 
indebtedness  in  the  preparation  of  this  manual  to  the  following 
works  on  analytical  chemistry:  H.  P.  Talbot's  Quantitative 
Chemical  Analysis;  J.  W.  Mellor's  Treatise  on  Quantitative  In- 
organic Analysis;  W.  F.  Hillebrand's  The  Analysis  of  Silicate 
and  Carbonate  Rocks;  F.  P.  Treadwell's  Lehrbuch  der  analytischen 
Chemie;  W.  C.  Blasdale's  Principles  of  Quantitative  Analysis; 
and  A.  Fischer's  Elektroanalytische  Schnellmethoden. 

G.  McP.  SMITH. 
UNIVERSITY  OP  ILLINOIS 
1918 


CONTENTS 
PART  I 

PAGE 

A.  INTRODUCTION       .        .        .        ; i 

Gravimetric  and  Volumetric  Analysis. 

B.  GENERAL  REMARKS  CONCERNING  QUANTITATIVE  WORK       .        .       3 

Neatness ;  Accuracy  and  Integrity ;  Economy  of  Time ;  Note- 
books; Reagents. 

C.  THE  OPERATIONS  or  ANALYTICAL  CHEMISTRY 

I.  Weighing -        7 

The  Balance ;  The  Use  and  Care  of  the  Analytical  Balance ; 
Determination  of  the  Zero-point;  Methods  of  Weighing; 
The  Calibration  of  a  Set  of  Weights;  Errors  Due  to  In- 
equalities in  Length  in  the  Beam  Arms ;  Errors  Due  to  the 
Buoyancy  of  the  Atmosphere. 

II.  Precipitation 20 

Qualities  Desirable  in  Precipitates  Which  Are  to  Be  Used  in 
Gravimetric  Determinations;  Colloidal  and  Fine-grained 
Precipitates;  The  Contamination  of  Precipitates;  The 
Theory  of  Precipitation. 

III.  Filtration  and  the  Washing  of  Precipitates     .        •.        .        .      30 

The  Selection  and  Use  of  Paper  Filters;  Wash  Bottles; 
Gooch's  Filtration  Crucible;  The  Theory  of  Washing 
Precipitates. 

IV.  The  Drying  and  Ignition  of  Precipitates         .        ...      37 

Drying  Ovens;  Desiccators;  Crucibles. 

V.  The  Evaporation  of  Liquids    .        ....        «        .      41 

VI.  The  Volumetric  Measurement  of  Liquids        .        .        .        -43 

Volumetric  Apparatus ;  Necessary  Precautions  in  the  Use  of 

Volumetric  Apparatus;    The   Calibration  of  Volumetric 

Apparatus. 

D.  THE  PREPARATION  OF  SAMPLES  FOR  ANALYSIS  .       .       .       .      51 

vii 


viii  CONTENTS 

PART   II 
GRAVIMETRIC  ANALYSIS 

PAGE 

EXERCISES  WITH  THE  BALANCE         .       .   .    ..-.    .        .       .        .      53 
THE  DETERMINATION  or  CHLORINE  IN  A  SOLUBLE  CHLORIDE          .      54 

THE  DETERMINATION  or  IRON  AND  OF  SULPHUR  IN  A  SOLUBLE  SUL- 
PHATE OF  IRON    .        .        ...      .      .",        .        .        ...      59 

THE  DETERMINATION  OF  SULPHUR  IN  AN  ORE         v        .        .        .  65 
THE  DETERMINATION  OF  PHOSPHORIC  ANHYDRIDE  IN  PHOSPHATE  ROCK  66 
THE  DETERMINATION  OF  CALCIUM  AND  MAGNESIUM  OXIDES  IN  LIME- 
STONE .        .        .        .        .        *        .        ...        .        .  70 

THE  DETERMINATION  OF  CARBON  DIOXIDE  IN  LIMESTONE       .        .  76 

THE  DETERMINATION  OF  SILICA  IN  A  REFRACTORY  SILICATE   .        .  So 

THE  DETERMINATION  OF  POTASH  IN  SOLUBLE  SALTS    '.  .       .        .  84 

THE  ELECTROLYTIC  DETERMINATION  OF  COPPER       ....  86 

PART   III 
VOLUMETRIC  ANALYSIS 

GENERAL  DISCUSSION        .       .     .  .       .  *     .....       .      97 

Fundamental  Principles ;  Reactions  Suitable  for  Volumetric  Pro- 
cesses; Determination  of  the  End-point;  General  Theory  of 
Indicators ;  The  Advantages  of  the  Volumetric  System ;  General 
Directions. 

A.  NEUTRALIZATION  METHODS:   ALKALIMETRY  AND  ACIDIMETRY      .     105 
Standard  Acid  Solutions ;  Standard  Alkali  Solutions ;  Indicators 
for  Use  hi  Alkalimetry  and  Acidimetry. 
THE  PREPARATION  AND   STANDARDIZATION   OF  APPROXIMATELY 

HALF-NORMAL  HYDROCHLORIC  ACID  AND  SODIUM  HYDROXIDE    109 
THE  DETERMINATION  OF  THE  TOTAL  ALKALINE  VALUE  OF  SODA 

ASH        .        .      ,,.• 113 

THE  DETERMINATION  OF  THE  NEUTRALIZATION  VALUE  OF  AN  Aero    1 1 5 
THE  DETERMINATION  OF  PROTEIN  NITROGEN  BY  THE  KJELDAHL 

METHOD  .    116 


CONTENTS  ix 

PAGE 

B.  METHODS  OF  OXIDATION  AND  REDUCTION    .        *       .       .        .    119 

Standard  Solutions ;  Indicators. 

1.  BICHROMATE  PROCESSES         .  ^    .        .....        .    120 

Fundamental  Principles. 

THE  PREPARATION  AND  STANDARDIZATION  or  APPROXIMATELY 
TENTH-NORMAL  BICHROMATE  AND  FERROUS  IRON  SOLU- 
TIONS .  .  .  .  .  i 121 

THE  DETERMINATION  OF  IRON  IN  SIDERITE  .  .  .  .124 
THE  DETERMINATION  OF  CHROMIUM  IN  CHROME  IRON  ORE  .  125 

2.  PERMANGANATE  PROCESSES 127 

Fundamental  Principles. 

THE  PREPARATION  AND  STANDARDIZATION  OF  AN  APPROXI- 
MATELY TENTH-NORMAL  SOLUTION  OF  POTASSIUM  PER- 
MANGANATE .'.:...'.- 128 

THE  DETERMINATION  OF  IRON  IN  HEMATITE  .  .  .130 
THE  DETERMINATION  OF  CALCIUM  IN  LIMESTONE  .  .  .133 
THE  DETERMINATION  OF  THE  OXIDIZING  VALUE  OF  PYRO- 

LUSITE 134 

THE  DETERMINATION  OF  PHOSPHORUS  IN  STEEL  .  .  .135 
THE  DETERMINATION  OF  MANGANESE  IN  AN  ORE  .  .  .139 

3.  IODOMETRIC  PROCESSES 142 

Fundamental  Considerations. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROXIMATELY 
TENTH-NORMAL  SOLUTIONS  OF  IODINE  AND  SODIUM  Tmo- 
SULPHATE         .        .....        •        •        •        •     J46 

THE  DETERMINATION  OF  ANTIMONY  IN  STTBNITE   .        .        .148 
THE  DETERMINATION  OF  LEAD  IN  AN  ORE    .        .        .        .150 

THE  DETERMINATION  OF  COPPER  IN  AN  ORE        .        .        .152 

C.  PRECIPITATION  METHODS       .        .        ...        .        .        .        .     155 

General  Discussion. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROXIMATELY 
TENTH-NORMAL  SOLUTIONS  OF  SILVER  NITRATE  AND  POTAS- 
SIUM THIOCYANATE  .  .,'...  .  .  .  156 

THE  DETERMINATION  OF  CHLORINE  IN  A  SOLUBLE  CHLORIDE    .    158 


X  CONTENTS 

PART   IV 

STOICHIOMETRY 

PAGE 

PRELIMINARY  DISCUSSION:   THE  SOLUTION  OF  TYPICAL  PROBLEMS      159 
PROBLEMS   .        .       .       .        .        ,       ,        .       ...        .166 

PART  V 

QUESTIONS          .       .       .       .       .       .       .       .  .       .    178 

APPENDIX 

PREPARATION  OF  THE  REAGENTS 193 

SULPHURIC  ACID-DICHROMATE  CLEANING  SOLUTION         .        .        .196 
ANALYTICAL  SAMPLES  FOR  THE  USE  OF  STUDENTS    .        .        .        .196 

APPARATUS  IN  THE  STUDENT'S  DESK 197 

LOGARITHMS       .       .       .." . 198 

ANTILOGARITHMS        .  200 

INTERNATIONAL  ATOMIC  WEIGHTS,  1917    .        .        .       Back  Cover  Sheet 


AN  INTRODUCTORY   COURSE  IN 

QUANTITATIVE   CHEMICAL   ANALYSIS 


PART    I 

INTRODUCTION 
A.   GRAVIMETRIC  AND   VOLUMETRIC  ANALYSIS 

QUANTITATIVE  analysis  has  for  its  object  the  determination  of 
the  quantities  of  the  elements  or  compounds  which  are  present 
in  particular  samples  of  material.  The  results  are  usually  ex- 
pressed in  terms  of  percentage,  ordinarily  by  weight ;  but  some- 
times, as  in  the  analysis  of  gases,  by  volume. 

The  procedure  to  be  employed  in  a  specific  case  will  often  de- 
pend upon  the  qualitative  composition  of  the  sample.  A  quali- 
tative analysis,  therefore,  should  always  precede  a  quantitative, 
unless  the  composition  of  the  sample  is  sufficiently  well  known. 

In  the  performance  of  quantitative  determinations  there  are 
two  principal  methods  of  procedure,  according  to  which  the  sub- 
ject is  subdivided  into  gravimetric  and  volumetric  analysis.  In 
addition,  there  are  gasometric  methods,  and  various  physical 
methods,  of  analysis ;  but  these  will  not  be  described  in  this  book. 

In  a  gravimetric  analysis,  a  weighed  sample  is  taken,  and  the 
substances  to  be  determined  are  separated,  one  after  another, 
either  in  the  free  state,  or  in  the  form  of  suitable  compounds. 
Each  final  product  is  weighed,  and,  from  its  weight,  the  weight, 
and  therefore  the  percentage,  of  the  corresponding  substance  in 
the  sample  can  be  calculated. 

The  substance  to  be  weighed  is  in  most  cases  separated  from 
solution  by  precipitation,  though  in  many  instances  it  is  deposited 
upon  a  weighed  cathode  or  anode  by  electrolysis.  Sometimes 
it  is  separated  from  other  substances  by  extraction  with  a  solvent, 
and  sometimes  in  the  form  of  a  gas,  the  weight  of  the  gas  being 
determined  either  by  absorbing  it  in  a  weighed  quantity  of  some 
substance  and  noting  the  increase,  or  by  noting  the  decrease  in 
weight  due  to  the  removal  of  the  gas  alone. 


IN  QUANTITATIVE  ANALYSIS 


In  a  volumetric  analysis  a  weighed  sample  is  also  taken,  but 
the  quantity  of  the  substance  to  be  determined  is  arrived  at 
by  causing  some  well-defined  reaction  to  take  place,  the  reagent 
being  added  from  a  burette,  in  the  form  of  a  solution  of  known 
concentration.  This  operation  is  called  titration.  From  the 
volume  of  the  solution  added,  it  is  easy  to  calculate  the  weight 
of  the  substance  present  in  the  sample. 

In  many  instances,  it  is  necessary  in  volumetric  analysis  also 
to  separate  the  substance  to  be  determined  from  interfering 
substances  present  with  it  in  the  sample  ;  but,  instead  of  mak- 
ing a  final  weighing,  the  substance  is  again  brought  into  solu- 
tion, in  suitable  form,  and  its  quantity  estimated  by  titration. 

In  order  to  illustrate  the  two  methods,  let  us  consider  the 
determination  of  chlorine  in  sodium  chloride. 

(a)  Gravimetric  Method.    The  weighed  sample  is  dissolved 
in  water,  the  solution  acidified  with  nitric  acid,  and  the  chlorine 
converted  into  insoluble  silver  chloride  by  means  of  an  excess  of 
silver  nitrate  solution.     The  precipitate  is  filtered  off,  washed, 
dried,  and  weighed.     From  its  weight,  the  weight  of  chlorine 
may  be  calculated,  as  follows: 

Cl 

-  X  wt.  of  precipitate  =  wt.  of  chlorine. 

And,  of  course,  wt  °f  chlori"e  Xioo  =  %  of  chlorine. 
wt.  of  sample 

(b)  Volumetric  Method.    The  weighed  sample  is  dissolved 
in  water,  the  solution  acidified  with  nitric  acid,  and  the  chlorine 
converted  into  silver  chloride  by  the  gradual  addition,  from  a 
burette,  of  a  silver  nitrate  solution  of  known  concentration. 
As  soon  as,  after  stirring  each  time  and  allowing  the  precipitate 
to  settle,  the  first  drop  is  added  which  fails  to  produce  a  pre- 
cipitate, the  reaction  is  known  to  be  complete  ;  and  the  number 
of  cubic  centimeters  required,  multiplied  by  the  chlorine  equiva- 
lent of  the  silver  nitrate  contained  in  each  cubic  centimeter, 
gives  directly  the  weight  of  chlorine  in  the  sample. 


INTRODUCTION  3 

B.   GENERAL  REMARKS   CONCERNING 
QUANTITATIVE  WORK 

Neatness.  The  drawers  and*  cupboards  of  the  desk,  and  all 
apparatus,  should  at  all  times  be;  neat  and  clean.  A  sponge  or 
an  old  towel  should  always  be  at  hand,  and  the  desk  top  and 
filter-stands  should  be  kept  dry  and  clean.  Vessels  should 
be  scrupulously  clean,  inside  and  out,  and  the  outer  surfaces  of 
beakers,  flasks,  etc.  should  be  wiped  dry  with  a  clean,  lintless 
towel,  before  use. 

If  the  inner  surfaces  of  funnels,  flasks,  etc.  become  contam- 
inated with  a  film  of  grease,  they  should  be  rinsed  either  with  a 
strong  solution  of  sodium  hydroxide  or  with  sulphuric  acid- 
dichromate  cleaning  solution.  (The  latter  may  be  prepared 
by  pouring,  cautiously  and  with  stirring,  4  volumes  of  con- 
centrated sulphuric  acid  (sp.  gr.,  1.84)  into  3  volumes  of  cold 
water,  and  saturating  the  resulting  hot  solution,  without  further 
heating,  with  powdered  sodium  or  potassium  dichromate.)  In 
extreme  cases  it  may  be  necessary  to  allow  the  apparatus  to 
stand  overnight  in  contact  with  this  solution. 

Accuracy  and  Integrity.  It  is  of  fundamental  importance  in 
quantitative  work  to  guard  against  loss  of  material  or  the  in- 
troduction of  foreign  matter.  All  filters  and  solutions  should 
be  kept  covered  to  protect  them  from  dust,  and  in  dissolving 
substances  for  analysis,  the  vessels  should  always  be  kept  covered 
to  prevent  mechanical  losses. 

Success  in  quantitative  work  demands  first  of  all  a  certain 
amount  of  dexterity  in  the  performance  of  the  mechanical  opera- 
tions involved.  Certain  individuals  are  able  to  acquire  this 
skill  with  comparative  ease,  but  the  majority  of  persons  can 
acquire  it  only  through  patient  and  persistent  application.  If 
the  student  finds  himself  unable  to  do  as  good  work  as  his  more 
experienced  or  more  fortunate  neighbor,  he  should  rather  devote 
his  energies  to  increasing  his  proficiency  than  to  trying  to  con- 
ceal his  lack  of  it.  Nothing  less  than  absolute  integrity  can  be 


4     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

demanded  of  an  analytical  chemist,  and  any  disregard  of  this 
principle  is  certain  to  be  fatal  to  his  success. 

Economy  of  Time.  An  economical  use  of  laboratory  hours 
is  best  secured  by  acquiring  a  thorough  knowledge  of  the  char- 
acter of  the  work  to  be  done  before  undertaking  it,  and  then 
arranging  the  work  so  that  no  time  shall  be  wasted  during  the 
evaporation  of  liquids  and  other  time-consuming  operations. 

At  least  two  determinations  should  be  in  progress  at  the  same 
time,  and  confusion  should  be  carefully  guarded  against  by  a 
free  use  of  labels.  In  general,  economy  of  time  results  from  the 
evaporation  or  filtration  of  several  solutions  at  once;  four  or 
more  precipitates  may  often  be  washed  in  the  time  required  for 
any  one  of  them,  if  taken  alone. 

Notebooks.  Notebooks  should  contain,  besides  the  analytical 
data,  descriptive  notes  regarding  any  special  difficulties  en- 
countered in  the  analysis  and  the  remedies  applied,  and  also 
any  incidents  in  the  course  of  the  analysis  which  might  influence 
the  results  injuriously. 

All  analytical  data,  such  as  records  of  weights  and  volumes, 
should  be  placed  upon  the  right-hand  page,  while  the  left-hand 
page  should  be  reserved  for  the  descriptive  notes,  the  calcula- 
tion of  factors,  of  the  amounts  of  reagents  required,  etc. 

All  analyses  should^be  made  in  duplicate,  and  in  general  a 
close  agreement  in  results  should  be  expected.  It  should,  how- 
ever, be  realized  that  "  check  results  "  do  not  furnish  con- 
clusive evidence  of  accuracy.  Since  check  results  depend  almost 
entirely  upon  the  prevalence  of  identical  conditions  throughout 
the  course  of  the  two  analyses,  they  are  apt  to  be  obtained  even 
when  inaccurate  methods  of  analysis  are  employed.  A  common 
fallacy  is  to  the  effect  that  no  part  of  the  work  need  be  performed 
more  carefully  than  that  part  which  is  necessarily  least  accurate. 
For  example,  it  is  said  that  if  a  certain  step  in  a  process  involves 
an  unavoidable  error  of  0.1%,  it  is  a  waste  of  time  to  attempt 
to  avoid  errors  in  other  parts  of  the  work  amounting  to  0.05% 
or  even  0.09%.  This  unfortunate  attitude  would  lead  to  the 


INTRODUCTION  5 

conclusion  that  if  a  method  cannot  yield  results  involving  an 
error  of  less  than  0.10%,  it  should  be  given  a  chance  to  depart 
o.io%-hwxo.09%,  if  there  are  n  other  places  where  errors  may 
occur.  Of  course  these  errors  may,  to  a  certain  extent,  counter- 
act one  another  in  effect,  but  there  is  no  assurance  that  they  will 
do  so.  Nevertheless,  when  a  certain  minimum  error  is  unavoid- 
able, e.g.  0.20%,  it  is  not  wise  to  expend  an  undue  amount  of 
time  in  trying  to  prevent  other  possible  errors  when  the  ratio  of 
these  errors  to  the  larger  error  is  very  small ;  because  in  such  a 
method  the  percentage  result  has  no  significance  whatever  beyond 
the  first  decimal.  It  is  a  good  rule  always  to  report  one  decimal 
place  further  than  the  one  that  is  considered  to  be  certainly  correct. 
All  records  should  be  dated,  and  all  observations  should  be 
recorded  at  once  in  the  notebook.  Records  should  never  be 
made  upon  loose  sheets  of  paper.  Since  the  neat  and  systematic 
arrangement  of  the  analytical  data  in  the  notebook  is  a  matter 
of  the  first  importance,  the  following  sample  right-hand  page  is 
given  as  a  suggestion  of  the  manner  in  which  such  records  should 
be  kept.  In  the  analysis  here  given,  it  is  uncertain  whether 
the  first  figure  after  the  decimal  should  be  4,  3,  or  2 ;  of  course, 
then,  nothing  is  known  concerning  the  value  of  the  second  figure, 
which  therefore  is  not  included  in  the  mean  result. 

DETERMINATION  OF  CHLORINE  IN  A  SOLUBLE  CHLORIDE 

I 

Sample  Tube,  etc.  8.4237 

Tube  minus  Sample  8.2377 

Wt.  of  Sample  0.1860 

Wt.  of  Crucible  5-3588 

Crucible+AgCl,  ist  time  5.7830 

2d  time  5.7828 

Wt.  of  Crucible  5.3588 

Wt.  of  AgCl  0.4240 

Per  cent  of  Chlorine  56.40 

Mean  Value =56.3% 


6     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Reagents.  Probably  the  greatest  hindrance  to  good  work 
in  otherwise  well-equipped  laboratories  is  the  difficulty  of  secur- 
ing satisfactory  reagents.  Also  much  of  the  glassware  on  the 
market  is  of  an  inferior  grade  and  utterly  unsuited  for  analytical 
work. 

The  habit  of  carefully  testing  reagents,  including  distilled 
water,  cannot  be  too  early  acquired ;  the  most  ceaseless  vigilance 
should  at  all  times  be  practiced  in  guarding  against  the  presence 
of  impurities  which  would  vitiate  the  analytical  work  under 
way.  As  is  generally  known,  a  "  C.  P."  label  is  no  guaranty 
whatever  of  the  purity  of  a  reagent,  and  the  "  guaranteed  "  or 
"  analyzed  "  reagents,  sold  at  high  prices,  are  at  times  worse 
than  products  for  which  no  claim  to  special  purity  has  been 
made. 

Acids  of  a  high  degree  of  purity  can  be  obtained  commer- 
cially, and,  although  exceptions  have  been  noted,  these  in  most 
cases  need  no  redistillation.  But,  owing  to  its  basic  nature, 
ammonia  ought  always  to  be  redistilled  at  short  intervals,  after 
first  shaking  it  up  with  slacked  lime  to  remove  any  carbonic 
acid.  Glass  stock  bottles  may  be  coated  inside  with  ceresin, 
to  prevent  contact  between  the  glass  and  the  ammoniacal 
solution. 

Owing  to  the  solvent  action  on  glass  of  many  solutions  of 
solid  reagents,  these  should  be  made  up  at  frequent  intervals 
in  limited  quantities,  or,  preferably,  the  solid  should  be  dis- 
solved as  wanted.  This  is  particularly  called  for  with  such 
reagents  as  ammonium  oxalate  and  microcosmic  salt,  and  alka- 
line "  magnesia  mixture  "  should  not  be  kept  in  contact  with 
glass. 

The  stopper  of  a  reagent  bottle  should  never  be  laid  upon 
the  desk,  but  should  always  be  held  in  the  fingers  until  returned 
to  the  bottle.  This  will  prevent  contamination,  whether  due 
to  an  interchange  of  stoppers,  or  to  some  other  cause.  The 
necks  and  mouths  of  such  bottles  should  of  course  be  kept 
scrupulously  clean. 


INTRODUCTION  7 

C.  THE  OPERATIONS  OF  ANALYTICAL  CHEMISTRY 

The  chief  operations  involved  in  analytical  work  which  can 
be  profitably  discussed  at  this  point  are  weighing,  precipitation, 
filtration,  and  the  washing  of  precipitates,  the  drying  and  igni- 
tion of  precipitates,  the  evaporation  of  liquids,  and  the  volumet- 
ric measurement  of  liquids. 

These  operations  will  be  described  in  the  following  sections, 
which  should  be  studied  carefully  by  the  beginner.  It  is  of 
prime  importance  for  success  as  an  analyst  to  pay  great  atten- 
tion to  details  and  scrupulously  to  avoid  any  conditions  which 
may  destroy  the  analysis,  or  lessen  confidence  in  the  accuracy 
of  the  data. 

The  adoption  of  the  suggestions  given  will  do  much  to  insure 
work  of  a  high  grade,  while  neglect  of  them  will  often  lead  to 
inaccurate  results  and  loss  of  time. 

I.   WEIGHING 

The  purpose  of  weighing  is  to  compare  the  quantity  of  matter 
in  a  specific  object  with  the  quantity  of  matter  in  a  given  stand- 
ard —  a  gram  or  kilogram  weight.  The  comparison  is  made 
on  the  balance  by  suspending  the  object  to  be  weighed  at  one 
end  of  a  beam,  and  the  weights  at  the  opposite  end  of  the  beam. 
The  beam  is  virtually  a  kind  of  lever,  and  the  mechanical  theory 
of  the  balance  is  founded  mainly  on  the  properties  of  levers. 

The  Balance.  The  beam  of  the  balance  is  supported  on  a 
central  knife-edge,  usually  of  agate,  which  rests  upon  a  plane 
agate  plate;  and  two  pans  for  supporting  the  masses  to  be 
compared  are  vertically  suspended  from  stirrups,  each  of  which 
has  an  agate  bearing  which  rests  on  a  knife-edge  fixed  at  one 
extremity  of  the  beam.  The  arms  of  the  balance  are  so  gradu- 
ated that  a  rider  (of  known  weight)  can  be  placed  on  the  beam  at 
any  required  distance  from  the  central  knife-edge. 

If  the  three  knife-edges  are  allowed  to  press  continually  upon 
their  agate  bearings,  they  soon  become  blunted,  and  wear  fur- 


8     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

rows  in  the  bearings.  In  order  to  prolong  the  life  of  the  knife- 
edges  and  bearings,  the  balance  is  provided  with  a  "  release  " 
which  separates  the  knife-edges  from  their  bearings  when  the 
balance  is  not  in  use.  If  the  balance  shows  signs  of  stiffness 
in  the  motion  of  beam  and  pans,  the  fault  should  be  investigated 
at  once.  The  defect  may  be  due  to  an  accumulation  of  dust 
between  the  knife-edges  and  their  bearings;  to  the  blunting  of 
the  knife-edges;  or  to  the  wearing  of  furrows  in  the  bearings. 
To  prevent  the  accumulation  of  dust,  and  also  to  prevent  the 
interference  of  air  currents  while  weighing,  the  balance  is  in- 
closed in  a  glass  case. 

In  order  to  render  small  movements  of  the  beam  perceptible, 
there  extends  downwards  from  its  center  a  long  pointer  which 
multiplies  the  rotational  displacement.  When  equilibrium  is 
established,  the  lower  end  of  the  pointer  should  come  to  rest  in 
front  of  the  zero  of  a  scale  which  is  located  immediately  behind 
this  end. 

The  conditions  which  must  be  satisfied  by  a  good  balance  are : 
(i)  The  balance  must  be  consistent.  It  must  give  the  same 
result  in  successive  weighings  of  the  same  body.  This  condition 
depends  upon  the  trueness  of  the  knife-edges.  (2)  The  balance 
must  be  accurate.  At  rest  the  beam  must  be  horizontal  when 
the  pans  are  empty,  and  when  equal  weights  are  placed  upon 
the  pans.  This  condition  depends  upon  the  equality  of  the 
two  arms.  (3)  The  balance  must  be  stable.  The  beam  after 
being  displaced  from  its  horizontal  position  must  return  to  its 
horizontal  position.  This  condition  depends  upon  the  adjust- 
ment of  the  center  of  gravity.  (4)  The  balance  must  be  sensi- 
tive. It  must  show  even  a  very  small  inequality  in  the  two 
masses  on  the  scale  pans.  This  condition  depends  largely  upon 
the  length  of  the  arms.  (5)  The  balance  must  oscillate  with 
reasonable  rapidity.  Short  beams  oscillate  more  rapidly  than 
long  ones. 

The  analytical  balance  will  perform  excellent  service  under 
the  proper  conditions,  but  great  care  in  its  use  is  essential  if  its 


INTRODUCTION  9 

accuracy  is  to  be  relied  upon.  It  should  be  located  in  a  room 
that  is  free  from  dust  and  fumes,  and  should  stand  upon  a  sup- 
port that  is  free  from  shocks  and  vibrations. 

The  Use  and  Care  of  the  Analytical  Balance.  The  following 
rules  embody  the  main  points  to  be  observed  in  the  use  and  care 
of  a  balance. 

(1)  Each  student  must  feel  a  personal  responsibility  for  the 
proper  use  of  his  balance;  carelessness  on  the  part  of  any  one 
is  apt  to  render  inaccurate  not  only  his  own  work,  but  also  that 
of  all  others  who  use  the  same  balance. 

(2)  The  balance  pans  should  be  brushed  off,  if  necessary,  and 
the  adjustment  of  the  balance  tested  before  use. 

The  balance  is  properly  adjusted  only  if  the  following  con- 
ditions are  fulfilled:  (a)  The  spirit  level  or  plumb  bob  inside 
the  balance  case  should  show  that  the  balance  is  level ;  (b)  the 
mechanism  for  raising  and  lowering  the  beam  should  work 
smoothly ;  (c)  the  pan  arrests  should  just  touch  the  pans  when 
the  beam  is  lowered;  (d)  the  pointer  should  rest  at  zero  when 
the  beam  is  raised,  and  also  when  it  is  lowered  so  that  the  pan 
arrests  touch  the  pans ;  and  (e)  the  pointer  should  swing  equal 
distances  on  either  side  of  the  zero-point  when  the  beam  is  set 
in  motion  without  any  load  on  the  pans.  In  the  latter  case,  if 
the  variation  does  not  exceed  two  divisions  on  the  scale,  it  is  hardly 
worth  while  to  disturb  the  balance  by  an  attempt  at  correction ; 
it  is  better  to  make  a  proper  allowance  for  the  small  zero  error. 

(3)  The  beginner  should  never  attempt  to  make  adjustments 
himself,  but  should  always  apply  to  the  instructor  in  charge. 

(4)  The  beam  should  never  be  set  in  motion  by  lowering  it 
upon  its  knife-edge,  nor  by  touching  the  pans,  but  rather  by 
means  of  the  rider ;  however,  there  is  a  "  trick  "  in  lowering  the 
pan  supports  so  that  the  oscillations  of  the  pointer  will  have  the 
required  amplitude. 

The  pans  should  be  arrested  and  the  beam  raised  before  any 
change  is  made  in  the  load  or  weights  on  the  pans  except  in  the 
case  of  the  small  fractional  weights,  when  it  is  only  necessary 


io     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

to  arrest  the  pans.  The  object  to  be  weighed  and  the  heavy 
weights  should  be  placed  in-  the  middle  of  their  respective  pans, 
since  a  heavy  load  near  the  edge  of  a  pan  is  apt  to  cause  trouble- 
some oscillations. 

The  beam  and  stirrups  should  never  be  left  upon  their  knife- 
edges,  and  the  motion  of  the  beam  should  be  arrested  only  by 
means  of  the  pan  arrests,  and  only  when  the  pointer  is  passing 
the  center  of  the  scale ;  otherwise  the  knife-edges  become  dull 
and  their  agate  bearings  furrowed. 

(5)  The  weights  should  be  cared  for  not  less  than  the  balance, 
and  should  be  standardized  by  the  analyst  unless  they  are  known 
to  be  in  satisfactory  condition. 

The  weights  should  be  handled  carefully,  and  only  with  the 
forceps  provided  for  that  purpose ;  they  should  never  be  touched 
with  the  fingers.  In  weighing,  the  weights  should  always  be 
placed  upon  the  same  pan,  and  they  should  be  taken  in  the  order 
in  which  they  occur  in  the  box,  the  larger  ones  first;  and  the 
weight  of  the  object  should  be  recorded  by  noting  the  vacant 
spaces  in  the  box.  The  record  so  obtained  should  be  checked 
as  the  weights  are  removed  from  the  pan.  In  this  way  errors 
are  not  likely  to  occur. 

(6)  No  analytical  sample  should  ever  be  placed  directly  upon 
the  balance  pan.     Furthermore,  the  object  to  be  weighed  should 
neither  be  warmer  nor  colder  than  the  air  in  the  balance  case. 
Currents  of  hot  air  may  impinge  on  the  arms  of  the  balance  and 
buoy  up  the  beam,  or  cause  one  arm  of  the  balance  to  expand 
unequally.1     If  the  object  is  colder  than  the  atmosphere  of  the 
balance  case,  moisture  may  condense  on  its  surface.     If  the 
body  to  be  weighed  is  likely  to  be  electrified  (e.g.  a  glass  weigh- 
ing tube),  it  should  be  allowed  to  stand  for  some  time  after  it 
has  been  wiped,  before  weighing. 

(7)  The  balance  case  should  be  closed  while  weighing  with 
the  rider,  so  as  to  avoid  currents  of  air. 

1  For  instance,  a  platinum  crucible  which  appeared  to  weigh  20.649  g.  when 
warm,  weighed  20.6920  g.  when  cold. 


INTRODUCTION  n 

As  soon  as  the  object  is  apparently  balanced  by  the  weights, 
the  beam  should  be  raised  and  again  lowered  into  place,  and  the 
observation  repeated.  This  will  assure  the  proper  alignment  of 
the  beam  and  pans  at  the  time  when  this  is  most  important. 

(8)  In  using  weighing  bottles  or  tubes,  the  vessel  should  be 
weighed  together  with  its  contents.     A  quantity  suitable  for 
analysis  should  then  be  "removed  without  loss,  and  the  vessel 
and  contents  again  weighed.     The  difference  in  weight  indicates 
the  quantity  of  sample  taken. 

Cork  stoppers  in  weighing  tubes  are  apt  to  change  in  weight, 
owing  to  varying  amounts  of  moisture  absorbed  from  the  at- 
mosphere. It  is  therefore  necessary,  before  weighing  out  a 
new  sample  from  it,  to  confirm  the  recorded  weight  of  a  tube 
which  has  been  unused  for  some  time. 

(9)  Errors  in  weighing  should  fall  well  within  the  limits  of  the 
experimental  error  due  to  the  analytical  operations.     If,  for  ex- 
ample, an  error  of  o.ooi  g.  were  made  in  weighing  out  a  gram 
sample  of  clay  containing  0.20%  of  MgO,  the  resulting  error 
in  the  determination  of  the  magnesia  could  be  no  greater  than 
0.1%  of  its  value ;  the  final  result  could  not  be  affected  by  more 
than  0.1%  of  0.20%,  i.e.  0.0002%.     This  is  negligibly  small. 
Suppose,  however,  that  an  even  smaller  error  of  0.0005  S-  were 
made  in  weighing  the  0.0055  g-  °f  Mg2P207 ;  this  would  repre- 
sent an  error  of  9%  of  the  magnesia  value,  which  is  inexcusable. 

(10)  If  any  substance  is  spilled  upon  the  pans,  or  if  anything 
at  all  appears  to  be  the  matter  with  a  balance,  the  fact  should 
at  once  be  reported  to  the  instructor  in  charge.     In  most  in- 
stances serious  injury  can  be  averted  by  prompt  action. 

Determination  of  the  Zero-point.  Lower  the  beam  and 
stirrups  upon  the  knife-edges  by  slowly  turning  to  the  left  the 
milled  head  at  the  front  of  the  balance  case.  Then  release  the 
pan  supports  by  gently  pressing  inwards  the  small  button,  also 
at  the  front  of  the  case,  and  with  the  beam  swinging  smoothly, 
make  a  consecutive  record  of  the  number  of  scale  divisions 
traversed  by  the  pointer  on  either  side  of  the  center.  Record 


12     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 


the  swings  to  the  left  as  negative  numbers  and  those  to  the  right 
as  positive  numbers,  making  four  observations  on  one  side  and 
three  on  the  other.  Take  the  average  of  each  column,  add  these 
averages  algebraically,  and  divide  the  sum  by  two.  The  quotient 
is  the  zero-point  of  the  balance,  i.e.  the  position  on  the  scale  at 
which  the  pointer  would  finally  come  to  rest. 
Example  : 


LETT 

RIGHT 

-6.8 
-6.6 
-6.4 
-6.3 

+4.7 
+4-5 
+4-3 

Average:                  —6.5 

Average  :                +4.5 

Zero-point  =  —  i.o. 

Two  methods  of  procedure  are  now  open  to  the  operator. 
He  may  either  make  his  weighings  with  reference  to  this  ob- 
served zero-point  or  he  may  adjust  the  balance  so  that  the  ob- 
served zero-point  is  the  actual  zero  of  the  scale.  The  first 
method  is  preferable,  unless  the  zero-point  is  more  in  error  than 
one  scale  division.  The  zero-point  is  apt  to  change,  and  it  must 
be  determined  each  day,  or  even  more  often. 

Methods  of  Weighing.  Weighings  smaller  than  0.005  g-  (°r 
o.oi  g.)  are  made  with  the  rider.  When  the  arms  are  divided  into 
five  divisions,  a  5-milligram  rider  is  used;  in  general,  the  rider 
should  weigh  as  many  milligrams  as  there  are  large  divisions  on  the 
beam  between  the  central  knife-edge  and  the  right-hand  stirrup 
support.  Each  division  on  the  beam  then  corresponds  to  a 
milligram. 

Ordinary  Method.  The  object  to  be  weighed  is  placed  upon  the 
left-hand  pan  of  the  balance  and  weights  upon  the  right-hand 
pan,  until,  finally,  the  further  addition  of  5  mg.  (or  10  mg.)more 
than  counterbalances  the  object.  This  weight  is  then  removed, 
the  balance  case  closed,  and  the  rider  adjusted  so  that  the  pointer 
swings  equal  distances  on  either  side  of  the  zero-point.  This 


INTRODUCTION  13 

method  of  weighing  is  very  common,  and  it  is  sufficiently  accurate 
for  ordinary  analytical  work.  If  necessary,  the  zero-point  of  the 
unloaded  balance  should  be  determined  before  each  weighing. 

In  special  cases,  as  in  the  calibration  of  a  set  of  weights,  it 
is  important  to  make  more  accurate  weighings.  It  is  here  best 
to  use  the  method  of  weighing  by  double  vibrations,  which  from 
the  following  description  may  appear  somewhat  laborious; 
but  the  labor  is  more  apparent  than  real. 

Method  of  Weighing  by  Double  Vibrations,  (a)  Find  the 
zero-position  of  the  pointer  in  the  case  of  the  unloaded  balance, 
according  to  the  method  already  described.  Let  us  suppose 
this  to  be  at  +0.1. 

(6)  Find  the  deviation  of  the  scale  per  milligram,  that  is,  the 
sensitivity  of  the  loaded  balance.  The  object  to  be  weighed  is 
placed  upon  the  left  pan,  the  weights  on  the  right  pan.  When 
the  weights  are  so  far  adjusted  that  an  additional  0.005  g-  (°r 
o.oi  g.)  would  be  too  much  (e.g.  weight  on  pan  =  11.216  g.), 
close  the  door  of  the  balance  case,  and  adjust  the  rider  until 
the  pointer  swings  on  both  sides  of  the  zero  of  the  scale.  Now 
find  the  position  of  rest,  which  we  will  suppose  to  be  at  +0.8. 
Move  the  rider  one  milligram  division  to  the  right,  and  again 
find  the  position  of  rest;  this  being  at,  say,  —2.1.  Hence,  the 
zero-point  is  displaced  -fo.8— (  — 2.i)  =  2.p  divisions  by  increas- 
ing the  weight  i  milligram;  or  2.9  scale  divisions  correspond  to 
i  milligram  for  the  given  load. 

(c)  Calculate  the  weight  of  the  load  on  the  pan.  From  the 
preceding  results,  it  follows  that  the  load  weighs  11.216+0;. 
The  zero-point  of  this  load  is  displaced  0.8—0.1=0.7  scale 
division.  Since  2.9  scale  divisions  correspond  to  i  milligram, 

0.7  scale  division  will  correspond  to  —=0.24  mg.     Hence  the 

2.9 

weight  of  the  body  is  11.216+0.00024  =  11.21624  g.  These 
calculations  may  be  summarized  in  the  formula 

Correction  =  + r  mg., 

a— o 


14     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

in  which  z  represents  the  zero-point  of  the  unloaded  balance; 
a,  the  zero-point  with  not  quite  enough  weight  on  the  right  pan  ; 
and  b,  the  zero-point  with  a  milligram  more  on  the  right  pan 
than  corresponds  to  a. 

Analytical  balances  will  rarely  indicate  with  certainty  less 
than  o.oooi  g.  Hence,  although  the  weight  may  be  calculated 
as  above  to  the  fifth  decimal,  it  should  generally  be  rounded  off 
by  dropping  the  fifth  decimal  and  raising  the  fourth  decimal  one 
unit  when  the  dropped  figure  exceeds  5. 

In  certain  cases,  as  in  the  calibration  of  volumetric  measuring 
apparatus,  it  is  necessary  for  the  weight  found  to  be  independent 
of  any  inequality  in  length  in  the  beam  arms.  In  such  cases, 
and  in  the  determination  of  absolute  weights  (reduction  to 
weights  in  vacua),  one  of  the  following  methods  should  be  used. 

Method  of  Gauss.  Weigh  the  object  first  in  one  pan,  then  in 
the  other.  Let  W  be  the  true  weight,  a  the  weights  required  to 
counterbalance  the  object  when  it  is  on  the  left  pan,  and  b  the 
weights  required  when  the  object  is  on  the  right  pan.  Accord- 
ing to  the  principle  of  moments  : 


=  ar 

That  is,  WHr  =  ablr^or  W2  =  ab  ; 

whence  W  —  Vab. 

Therefore  the  true  weight  is  the  square  root  of  the  product  of 
the  two  observed  weights. 

Borda's  Method  of  Weighing  by  Tares.  Here  the  object, 
placed  on  the  right-hand  pan,  is  balanced  by  a  suitable  tare 
(weights,  wire,  beaker  containing  shot,  etc.)  on  the  left-hand  pan. 
The  object  is  then  removed,  and  weights  are  added  in  its  place 
until  equilibrium  is  restored.  These  weights  are  necessarily 
the  same  in  value  as  the  object  for  which  they  substitute,  irre- 
spective of  differences  in  the  arms. 

The  Calibration  of  a  Set  of  Weights.  Fairly  accurate  weights 
can  be  purchased  for  a  reasonable  sum,  and  for  most  analytical 
work  the  inaccuracies  of  the  better  class  of  weights  are  negligibly 


INTRODUCTION  15 

small  in  comparison  with  the  errors  of  experiment,  and  the  im- 
perfections in  the  analytical  processes. 

An  analyst,  however,  should  know  that  his  weights  are  suffi- 
ciently accurate,  and  for  this  reason  he  should  calibrate  the 
weights.  The  errors  due  to  imperfections  in  the  weights  can 
easily  be  reduced  to  o.oooi  g.  The  weights  should  be  tested 
at  periodic  intervals,  say  once  or  twice  a  year,  depending  upon 
the  frequency  with  which  they  are  used. 

In  special  cases,  e.g.  in  the  calibration  of  volumetric  apparatus, 
absolute  weights  may  be  required,  but  for  general  analytical 
work  absolute  weights  are  not  necessary.  If  the  weights  are 
consistent  with  one  another,  their  absolute  values  have  no  in- 
fluence upon  the  accuracy  of  an  analysis. 

Before  beginning  the  calibration,  distinguish  all  separate 
pieces  of  the  same  denomination  by  marking  them  with  a  small 
prick  punch.  One  of  the  two  lo-gram  pieces  may  be  marked  ('), 
two  of  the  i -gram  pieces  (')  and  ("),  etc. 

The  method  of  weighing  to  be  followed  in  the  calibration 
will  depend  upon  the  degree  of  equality  in  the  lengths  of  the 
beam  arms.  If  they  are  unequal,  either  the  method  of  Gauss 
or  that  of  Borda  may  be  used  for  comparing  the  weights.  If  the 
method  of  Borda  is  used,  a  second  set  of  weights  will  be  found 
convenient  for  the  substitutions.  This  method  involves  less 
work  in  calculating  then  does  that  of  Gauss.  If  the  beam  arms 
are  essentially  equal,  the  simple  method  of  double  vibrations  is 
used,  without  the  necessity  of  a  correction.  In  any  case,  the 
following  comparisons  are  made,  with  the  use  of  the  rider  to 
obtain  equilibrium. 

I 

GRAM  WEIGHTS 

i  against  i' 

1  against  i" 

2  against  i  +  i' 

5  against  2-fi-fi'+i" 


16     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

10  against  5+2+1  +  1'+!" 
10  against  10' 
20  against  10+10' 
Etc. 

II 

FRACTIONAL  WEIGHTS 

500  against  200+ 100+100'+  50+  20+10+10' +5+rider  at  5 
200  against  100+ 100' 
100  against  ioo/l 

100  against  so+2o+io+io'+5+rider  at  5 
50  against  2o+io+io'+5+rider  at  5 
20  against  10+ 10' 
10  against  10'. 
10  against  5+ rider  at  5 
5  against  rider  at  5 

Also  unmarked  i-gram  piece  against  all  of  the  fractional  pieces + rider 
at  5 

The  calculation  of  the  weight  of  each  piece  of  i  gram  and 
upward  is  made  upon  the  arbitrary  assumption  that  the  un- 
marked i -gram  weight  is  correct.  In  calculating  the  weights 
of  the  fractional  pieces,  we  first  assume  the  unmarked  lo-mg. 
piece  as  a  standard  and  calculate  provisional  weights  for  each 
of  the  other  fractional  pieces  upon  this  basis.  We  then  add 
these  provisionally  corrected  weights  and  determine  by  compar- 
ing the  results  with  their  collective  weight  as  found  in  terms  of 
the  standard  i-gram  piece,  how  much  each  weight  must  be 
further  corrected.  If,  for  example,  the  sum  of  the  provisional 
fractional  weights  were  found  to  be  5  (rider) +  5 +  10+10+ 19.9 
+49.7  +  100+100.1  +  200.1+499.2  =  999.0  mg.,  or  0.9990  g., 
while  their  collective  weight  in  terms  of  the  i-gram  standard 
=  i.ooi7g.,  then  each  of  the  provisional  values  should  be  multi- 
plied by  — =1.0027.  I*1  thig  way,  we  obtain  the  weight  of 

0.9990 

each  piece  in  the  set  in  terms  of  the  unmarked  i-gram  weight. 

If  desired,  we  can  then  find  the  exact  value  of  the  unmarked 
i-gram  piece  in  terms  of  an  absolute  standard  weight.  For 
example,  if  the  unmarked  i-gram  weight  is  found  by  comparison 


INTRODUCTION  17 

to  weigh  0.9998  g.,  we  have  simply  to  multiply  each  weight  in 
the  table,  based  upon  the  unmarked  i-gram  weight  as  a  standard, 
by  0.9998,  in  order  to  obtain  the  weight  of  each  separate  piece 
of  the  set  in  terms  of  the  absolute.  standard. 

Errors  Due  to  Inequalities  in  Length  in  the  Beam  Arms.  In 
the  preceding  discussion,  it  has  mainly  been  assumed  that  the  two 
arms  of  the  beam  are  equal  in  length.  This  is  not  really  the  case. 
It  is  mechanically  impossible  to  insure  perfect  equality.  To  find 
the  relative  lengths  of  the  arms,  place  (corrected;  weights  of  the 
same  nominal  value  —  say,  50  grams  —  upon  each  pan,  and  bring 
the  balance  into  equilibrium  by  means  of  the  rider.  Interchange 
the  weights  on  the  two  pans,  and  again  bring  the  balance  into 
equilibrium  by  means  of  the  rider.  Call  the  two  weights  W  and 
wr  and  let  /  and  r  respectively  denote  the  additional  weights  re- 
quired for  equilibrium  on  the  left  and  right  sides.  Then,  on  the 
first  weighing,  w+l  =  W  ;  and,  on  the  second  weighing,  W  =  w+r. 
Let  L  and  R  respectively  denote  the  length  of  the  left  and  right 
arm.  Then  from  the  law  of  levers, 

L(w+t)  =  RW;  zudLW  =  R(w+r) 

Solving  each  of  these  equations  for  W  ',  and  equating  the  results, 
we  find  that 


whence, 


Suppose,  for  example,  that  the  weighings  were  found  to  be : 

LEFT  RIGHT 

50  =  2o+io+io'+io"+o.i3  mg. 

Here,  then,  /=  —0.00013,  r  =  +0.00019,  and  ^  =  50  g.     Conse- 
quently if  in  the  above  expression  we  let  R  =  i,  we  have 

r        \w+r 

L  =  A I  — —  =  i .  000003  2 
\w+l 

i.e.  L :  R  =  i  .0000032  :  i 


i8     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

With  this  ratio,  —=1.0000032,  a  weight  w  on  the  left  pan  of 
K. 

the  balance  will  be  equivalent  to  a  weight  wX  1.000003 2  on  the 
right  pan.  Hence,  if  a  substance  on  the  left  pan  balances  the 
weight  50  g.  on  the  right  pan,  the  weight  of  the  substance  is 
50.0000X1.0000032  =  50.0016  g.  —  an  error  of  only  0.003  Per 
cent.  There  is  therefore  no  need  to  apply  the  correction.  Each 
balance  has  its  own  constant  L :  R  for  a  given  load ;  the  numer- 
ical value  of  the  ratio  varies  with  the  different  loads. 

Most  analytical  balances  do  not  require  a  correction  on  account 
of  inequalities  in  the  arms ;  the  arms  are  usually  made  sufficiently 
exact.  Furthermore,  if  the  weights  are  always  placed,  say,  on 
the  right  pan,  such  a  correction  is  unnecessary  in  ordinary 
analytical  work,  because,  although  there  may  be  differences 
between  the  apparent  and  true  weights  of  the  substances  weighed, 
these  differences  are  proportional  to  the  true  weights  and  there- 
fore do  not  affect  the  ratios  obtained. 

Errors  Due  to  the  Buoyancy  of  the  Atmosphere.  The  assump- 
tion is  made  that,  if  two  bodies  are  equal  in  weight  at  the  same 
tune  and  place,  they  contain  the  same  mass  or  quantity  of 
matter.  This  is  only  true  if  the  two  bodies  have  the  same 
volume,  or  if  the  weighing  is  carried  out  in  a  vacuum.  A  body 
weighed  in  air  is  buoyed  up  by  a  pressure  equivalent  to  the 
weight  of  the  air  which  it  displaces.  Suppose  that  exactly 
100  grams  of  platinum  (sp.  gr.  21.55)  are  weighed  in  air  with 
brass  weights  (sp.  gr.  8.4).  Then  4.5  cc.  of  air,  at  say  20°  and 
760  mm.,  i.e.  about  0.0054  g.,  are  displaced  by  the  platinum; 
while  the  weight  of  the  air  displaced  by  100  grams  of  brass  is 
0.0143  g.  Hence,  the  weight  of  brass  which  exactly  counterpoises 
100  grams  of  platinum  is  100+ (0.0143  —  0.0054)  =  100.0089  g. 
The  buoyancy  of  the  air  thus  produces  a  sensible  effect  whenever  the 
volume  of  the  load  differs  materially  from  the  volume  of  the  weights.1 

1  To  eliminate  the  buoyancy  correction  due  to  variations  in  temperature  and 
pressure  during  the  same  experiment,  it  is  customary,  in  weighing  bulky  glass 
apparatus  (potash  bulbs,  etc.),  to  use  a  similar  piece  of  apparatus  as  a  counterpoise. 


INTRODUCTION 


The  arithmetic  of  the  above  calculation  is  summarized  in  the 
formula : 

Corrected  weight  ±=w-\-co  ( ) 

\5       SiJ 

in  which  w  represents  the  apparent  weight  of  the  object ;  s,  the 
specific  gravity  of  the  object;  si,  the  specific  gravity  of  the 
weights ;  and  «,  the  weight  of  a  cubic  centimeter  of  air  under 
the  conditions  prevailing  at  the  time  of  the  experiment. 

To  illustrate  the  effect  of  the  buoyancy  of  air  on  the  different 
substances  usually  weighed  in  clay  analyses,  the  following  table 
may  be  quoted : 


ERROR  PER  GRAM  OF 

SUBSTANCE  WEIGHED 

SUBSTANCE  WEIGHED 

GRAVITY 

With  Brass 
Weights 

With  Platinum 
Weights 

Clay  . 

2.^ 

O.OOO3 

C.OCO4. 

Silica                

2.23 

0.0004 

o  0005 

Aluminum  oxide       .... 
Ferric  oxide     

3.85 
^.12 

O.OOO2 
O.OOOI 

0.0003 

O.OOO2 

Magnesium  pyrophosphate    . 
Calcium  oxide 

2.40 
2  QO 

0.0003 

o  0003 

O.OOO4 

o  0004 

Potassium  chloride  .... 
Sodium  chloride 

I.Q9 
2.13 

0.0004 

0.0004 

0.0006 
o  0005 

Potassium  chloroplatinate 

3-34 

O.OOO2 

0.0003 

In  ordinary  analytical  operations  we  have  to  deal  with  differ- 
ences in  weight,  and  with  ratios,  not  with  absolute  weights. 
When  the  amount  of  a  precipitate  is  determined  from  the  dif- 
ference in  the  weight  of  an  empty  crucible  and  of  the  crucible 
plus  the  precipitate,  the  buoyancy  correction  is  not  needed  for 
precipitates  with  a  specific  gravity  near  that  of  the  substance 
undergoing  analysis.  If,  however,  the  specific  gravities  are 
widely  separated,  it  may  be  worth  while  to  correct  for  buoyancy. 
For  instance,  since  the  specific  gravities  of  pyrites  and  barium 
sulphate  are  nearly  equal,  it  would  be  a  waste  of  time  to  correct 
for  buoyancy  in  determining  sulphur  in  a  sample  of  pyrites. 


20     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

On  the  other  hand,  in  standardizing  a  solution  of  silver  nitrate 
by  precipitating  silver  chloride  from  a  specific  weight  of  the 
solution,  the  buoyancy  of  air  may  affect  the  result  by  xtr  of 
one  per  cent. 

Summary.  The  foregoing  discussion  serves  to  show  that,  in 
work  involving  delicate  measurements,  it  is  always  advisable 
to  make  an  estimate  of  the  influence  of  the  various  sources  of 
error  on  the  final  result.  These  errors  can  only  be  neglected 
when  their  effect  is  small  in  comparison  with  the  error  derived 
from  other  sources.  The  chief  sources  of  error  commonly  intro- 
duced in  the  balance  room  are  those  arising  from :  (i)  varia- 
tions in  the  zero-point  of  the  balance ;  (2)  inconsistent  weights ; 
(3)  inequalities  in  length  in  the  beam  arms ;  and  (4)  the  buoy- 
ancy of  the  air.  In  weighings  making  any  pretense  to  "  accuracy 
to  the  TO  milligram/'  the  following  points  should  be  noted : 

(1)  The  zero-point  of  the  unloaded  balance  should  be  deter- 
mined and  made  use  of  in  each  weighing. 

(2)  The  weights  should  be  calibrated,  and  periodically  checked 
for  consistency  among  themselves. 

(3)  The  errors  due  to  inequality  in  length  in  the  beam  arms 
can  be  neglected  in  ordinary  analytical  work. 

(4)  The  correction  of  the  weighings  for  the  buoyancy  of  air 
is  necessary  when  the  determination  involves  the  weighing  of 
substances  with  appreciably  different  specific  gravities.1 

II.  PRECIPITATION 

Qualities  Desirable  in  Precipitates  Which  Are  to  be  Used  in 
Gravimetric  Determinations.  Precipitation  is  made  use  of 
more  often  than  any  other  means  for  the  separation  of  inorganic 
substances.  But,  in  carrying  out  such  separations,  precipitates 
should  conform  as  nearly  as  possible  to  the  following  ideal 
specifications:  (i)  The  precipitate  should  be  insoluble  in  the 

xln  general  analytical  work  this  correction  can  almost  always  be  neglected, 
since  the  resulting  error  is  usually  overshadowed  by  the  errors  associated  with  the 
preparation  of  the  precipitates  for  the  balance.  ..*+  •• 


INTRODUCTION  21 

mother  liquid,  and  also  in  the  wash  liquid  to  be  used ;  (2)  it 
should  be  compact,  easy  to  filter  and  wash ;  (3)  it  should  be  a 
pure  chemical  substance  of  known  percentage  composition; 
and  (4)  it  should  either  be  stable  and  non-volatile  on  heating, 
or  it  should  yield  upon  ignition  a  pure,  non-volatile  substance 
of  known  composition.  The  last  two  conditions  are  of  especial 
importance  if  the  precipitate  is  the  substance  finally  to  be 
weighed.  Moreover,  other  things  being  equal,  it  is  conducive 
to  accuracy  if  a  precipitate  can  be  obtained  which  contains  a 
low  percentage  of  the  substance  under  investigation  (cf .  Part  IV, 
Problem  46). 

But  few  processes  satisfy  all  these  requirements,  and  in  the 
case  of  any  analytical  process  it  is  important  to  know  what 
conditions  favor  and  what  conditions  hinder  the  separation  and 
purification  of  a  given  precipitate.  There  are  a  few  general 
principles  of  such  wide  applicability  that  they  should  be  con- 
stantly borne  in  mind.  Their  discussion  follows. 

Colloidal  and  Fine-grained  Precipitates.  Finely  divided 
precipitates,  such  as  newly  precipitated  silver  chloride,  barium 
sulphate,  calcium  oxalate,  etc.,  are  particularly  liable  to  pass 
through  the  filter ;  furthermore,  they  tend  in  large  measure  to 
stop  up  the  pores  of  the  filter,  and  thus  to  increase  the  time 
required  for  filtration  and  washing.  Hence,  the  analyst 
employs  various  artifices  in  order  to  enlarge  the  size  of  the 
particles. 

(i)  The  grain  size  can  frequently  be  increased  by  allowing 
the  fine  grains  which  originally  separate  to  digest  in  the  pre- 
cipitation liquid.  This  change  is  more  rapid  in  the  hot,  than 
in  the  cold,  mother  liquid.  In  the  case  of  crystalline  substances, 
it  often  happens  that  the  finer  grains,  which  (owing  to  differences 
in  surface  tension)  are  somewhat  more  soluble  than  the  coarser 
ones,  redissolve;  and  since  the  solution  is  then  supersaturated 
in  respect  to  the  coarser  grains,  these  are  augmented  in  size  by 
the  surface  deposition  of  the  material  furnished  by  the  finer 
grains. 


22     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  boiling  of  liquids  containing  colloidal  substances  fre- 
quently leads  to  the  flocculation  of  a  large  number  of  fine  par- 
ticles into  a  smaller  number  of  coarser  aggregates. 

(2)  Precipitates  produced  in  hot  solutions  are  often  coarser- 
grained  than  if  produced  in  cold  solutions.     From  what  has 
just  been  said,  the  reasons  for  this  fact  will  be  apparent. 

(3)  The  flocculation  of  a  precipitate  which  separates  in  a 
colloidal  condition  is  frequently  caused  by  the  salts  present  in 
the  mother  liquid.     When  these  salts  have  been  almost  removed, 
during  the  washing,  the  colloidal  precipitate  is  apt  to  be  defloc- 
culated,  and  it  may  then  give  a  turbid  filtrate,  or  become  so 
slimy  as  to  be  almost  impermeable  to  the  wash  liquid.     In  such 
cases  it  is  necessary  to  wash  the  precipitate  either  with  boiling 
water,  or,  better,  with  the  solution  of  an  electrolyte  which  will 
prevent  the  deflocculation  of  the  precipitate,  and  which  can  be 
easily  removed  by  drying  or  ignition.     Sometimes  dilute  acids 
can  be  used,  but  usually,  for  obvious  reasons,  we  have  to  de- 
pend upon  volatile  ammonium  salts. 

The  Contamination  of  Precipitates.  Finally,  it  should  be 
noted  that  the  finer  the  grain  of  the  precipitate,  the  greater  will 
be  the  quantity  of  contaminating  salts  likely  to  be  retained  by 
the  wet  precipitate.  The  salts  appear  to  be  retained  by  a  kind 
of  surface  attraction,  called  adsorption,1  and,  since  fine-grained 
precipitates  expose  a  larger  surface  of  separation  between  the 
solid  and  the  liquid  phases,  and  also  because  of  their  compact- 
ness, the  fine  grained  precipitates  are  more  difficult  to  wash 
clean  than  those  of  coarser  texture.  Colloidal  gelatinous  pre- 
cipitates like  ferric  and  aluminum  hydroxides  are  in  an  ex- 
tremely fine  state  of  subdivision,  and,  in  consequence,  they  are 
most  difficult  to  wash  clean. 

In  addition  to  their  tendency  to  be  contaminated  by  the 
adsorption  of  salts,  precipitates  are  also  frequently  liable  to 
contamination,  owing  to  the  formation  during  precipitation  of 

1  But  see  "The  Contamination  of  Precipitates  in  Gravimetric  Analysis,"  G. 
McP.  Smith  :  Journal  of  the  American  Chemical  Society,  vol.  jp,  pp.  1152-73  (1917). 


INTRODUCTION  23 

more  or  less  stable  insoluble  complexes  (and,  in  rare  cases, 
possibly,  to  the  carrying  down  of  foreign  substances  by  the 
precipitate  in  a  state  of  solid  solution) .  These  impurities,  in  what- 
ever form  they  may  be  present,  cannot  be  completely  removed  by 
washing,  and  the  wash  water  will  frequently  fail  to  show  any  in- 
dication of  the  impurities  which  are  still  present  in  the  precipitate. 

It  is  therefore  often  advisable  to  redissolve  the  precipitate, 
and  to  repeat  the  precipitation.  The  objectionable  impurity 
divides  itself  in  a  more  or  less  definite  concentration  ratio  be- 
tween the  precipitate  and  the  mother  liquid.  A  relatively  large 
amount  may  be  retained  by  the  precipitate  in  the  first  precipita- 
tion, but  in  a  second  precipitation,  when  only  that  amount  of 
salt  retained  by  the  first  precipitate  is  in  solution,  the  division 
of  the  undesirable  substance  between  the  precipitate  and  the 
solution  in  the  given  concentration  ratio  means  that  a  much 
smaller  quantity  of  impurity  will  be  retained  by  the  second 
precipitate.  Repeated  precipitations  will,  in  general,  soon  carry 
the  amount  of  impurity  outside  the  range  of  the  balance ;  but, 
in  carrying  out  such  operations,  the  solubility  relations  of  the 
precipitate  itself  should  never  be  left  out  of  consideration. 

The  Theory  of  Precipitation.  Reversible  Reactions.  The 
reactions  which  are  made  use  of  in  analytical  chemistry  belong 
for  the  most  part  to  the  reversible  type.  Instead  of  running  to 
completion,  the  system  may  take  up  a  state  of  equilibrium  be- 
tween the  initial  stage  and  that  of  the  completed  reaction,  and 
a  certain  quantity  of  the  substance  under  investigation  is  apt 
to  escape  our  notice.  This  is  especially  true  in  many  reactions 
involving  precipitation,  neutralization,  oxidation,  etc. 

Furthermore,  many  reactions  are  more  or  less  influenced  by 
the  presence  of  certain  substances,  and  it  is  obvious  that  a 
thorough  knowledge  of  the  processes  which  take  place  in  such 
cases  will  be  of  the  greatest  service  to  the  analytical  chemist. 

Therefore,  it  is  of  primary  importance  in  analytical  chemistry 
to  study  each  process  thoroughly  in  detail,  with  the  object  of  finding 
out  and  understanding  the  conditions  which  will  be  most  favorable 


24     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

for  the  practical  completion  of  each  and  every  reaction  involved  in 
the  process.  ' 

Indispensable  guides  for  such  studies  are  found  in  the  ionic 
theory  and  the  law  of  mass  action.  It  is  taken  for  granted  that, 
at  this  point,  the  student  is  already  sufficiently  familiar  with 
the  qualitative  conception  of  ionization. 

Degree  of  Ionization.  In  a  dilute  aqueous  salt  solution,  the 
greater,1  and  by  far  the  most  active,  portion  of  the  salt  is  almost 
invariably  ionic.  But  with  acids  and  bases  there  is  a  wider 
range,  and  of  these  a  larger  number  are  less  highly  ionized ;  but 
even  here  the  ions  are  nearly  always  much  more  active  than  the 
non-ionized  molecules.  The  acids  and  bases  that  are  commonly 
called  "  strong "  are  highly  ionized,  i.e.  their  solutions  are 
especially  active  as  acids  or  bases  because  they  contain  high 
hydrogen,  or  hydroxide-ion,  concentrations. 

Composition  of  the  Ions.  It  is  usual  to  assume  the  simplest 
possible  compositions  for  the  ions  formed  upon  the  dissociation 
of  any  given  electrolyte.  A  more  careful  study  of  the  subject, 
however,  shows  that  the  ionization  of  even  simple  electrolytes 
may  be  a  very  complicated  process.  It  is  known,  for  example, 
that  sulphuric  acid  contains  ions  of  the  formula  HS04~,  in  ad- 
dition to  SO4 —  ions,  and  that  phosphoric  acid  yields  ions  of 

the  formulas  H2P04-,  HPO4— ,  and  PO4 .     All  of  these  are 

probably  more  or  less  highly  hydrated;  even  hydrogen  and 
hydroxide  ions  are  supposed  to  be  hydrated  in  aqueous  solu- 
tion. Further,  many  metallic  ions  show  a  decided  tendency  to 
exist  in  combination  with  certain  molecules  and  radicals,  as  OH2, 
OH,  NH3,  NH2,  CN,  C2O4,  P04,  Cl,  etc. ;  but  in  very  dilute 
solutions  these  complexes  are  apt  to  be  more  or  less  highly 
dissociated  into  their  constituents. 

The  Law  of  Mass  Action  as  Applied  to  Ionic  Equilibria.     In 
aqueous  solution,  acetic  acid  is  supposed  to  ionize  as  follows : 
HC2H302  Z£±  H++C2H302- 

1  Noteworthy  exceptions  are  mercuric  chloride  and  cyanide,  lead  acetate,  and 
a  few  others. 


INTRODUCTION 


The  quantity  of  the  molecular  acid  that  is  ionized  per  unit  of 
time  in  a  given  volume  of  the  solution  is  proportional  to  the 
concentration  of  the  non-ionized  molecules,  Cnc2H3o2,  while  the 
quantity  of  the  molecular  acid  that  is  simultaneously  formed 
by  the  union  of  the  ions  depends  upon  the  frequency  of  the  en- 
counters of  the  two  kinds  of  ions,  which  in  turn  is  proportional 
to  the  product  of  their  concentrations,  CH+xCc,Hs02-. 
The  speeds  of  the  respective  actions  will  therefore  be 

•5*1  =  CHc2H3o2  X  FI  and  .5*2  =  CH+  X  Cc2H8o2- 


in  which  FI  represents  the  intrinsic  tendency  of  HC2H3O2  to 
ionize,  and  FZ  that  of  H+  and  C^H-^Oz'  to  combine. 

When  equal  amounts  of  material  are  being  transformed  each 
way,  i.e.  at  equilibrium,  Si  =  S2,  and  therefore 

CnczHgOz  XFi  =  CH+  XCc2H,o,- 

or  CH+  xCc2H8o2-  _  FI  _  & 

--       —  ft- 


-^,  being  the  ratio  of  two  constants,  is  constant;  and  the  value, 
rz 

k,  of  this  ratio  of  the  affinities  driving  the  opposed  actions  is 
called  the  affinity  constant  of  the  reversible  reaction.  At  any 
given  temperature,  provided  the  solution  is  dilute,  the  numerical 
value  of  k  remains  the  same  no  matter  what  the  total  concentra- 
tion of  the  solution  may  be.1  In  the  case  of  acetic  acid,  for 
example,  the  following  figures  have  been  obtained,  at  18°,  from 
conductivity  determinations. 


TOTAL  MOLAL 
CONCENTRATION 
OF  AGED 

PROPORTION 
IONIZED 

MOLAL  CONCENTRATION 
OF  H+  AND  OF  AC- 
(Cn+  AND  Cc2H3o2-) 

MOLAL  CONCENTRATION 
OFHAc 

(CHC2H302) 

I.OOO 
O.IOOO 
0.0100 

0.0041 
0.0130 
0.0407 

O.OO4I 
O.OOI3O 

O.OOO407 

1.000-0.0041 
O.IOOO-O.OOI3O 
0.01000-0.000407 

1  When  data  such  as  the  following  are  applied  to  cases  of  soluble,  highly  ionized 
substances,  the  ^-values  so  obtained  for  any  given  compound  are  usually  far  from 
constant.  The  general  conclusions  arrived  at  through  the  application  of  such 
data  are,  however,  as  a  rule,  not  invalidated  by  this  fact. 


26     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 
Substituting  these  figures  in  equation  (i)  above,  we  get : 

(0.0041)2  ,      (o.oois)2 

-V^  =0.0000160 ;    —  =0.0000171 ; 

0.996  0.0987 

(0.000407) 2 

and — —  =0.0000172. 

0.00959 

It  is  seen  that,  although  the  third  solution  is  a  hundred  times 
more  dilute  than  the  first,  and  although  the  degree  of  ionization 
has  increased  tenfold,  the  value  of  k  is  the  same  in  both  cases. 

The  Common-Ion  Effect.  When,  through  the  presence  of 
two  substances  which  furnish  an  ion  in  common,  the  concentra- 
tions of  the  positive  and  negative  ions  of  an  ionogen  are  unequal, 
the  law  of  mass  action  still  holds. 

Let  us  imagine,  for  example,  that,  by  mixing  equal  volumes 
of  the  double-molal  solutions,  a  solution  is  obtained  which  is 
uni-molal  in  respect  to  acetic  acid  and  also  to  sodium  acetate. 
Let  us  further  suppose  that  the  equilibria  which  exist  in  the 
mixture  have  been  established  in  two  separate  stages,  as  follows : 
(i)  that  the  concentrations  of  each  undissociated  compound 
and  its  ions  have  changed  from  those  which  exist  in  a  double- 
molal,  to  those  which  exist  in  a  uni-molal  solution  of  that  com- 
pound; and  (2)  that  the  concentrations  of  all  the  substances 
present  have  changed  from  those  which  exist  in  the  separate 
uni-molal  solutions  of  the  compounds,  to  those  which  exist  in 
the  mixture  which  is  uni-molal  in  respect  to  each  compound. 
Let  us  now  consider  this  latter  stage  in  detail. 

In  uni-molal  solution,  sodium  acetate  is  0.53  ionized,  while 
acetic  acid  at  that  concentration  is  only  0.004  ionized.  Each  com- 
pound furnishes  acetate  ions,  and  the  acetate  ions  present  are 
all  available,  either  for  union  with  sodium  ions,  or  for  union  with 
hydrogen  ions.  Initially,  therefore,  in  the  case  of  the  sodium 

acetate,  we  have  a53X°'534  >klt  instead  of  °'53Xo.53  =fa 
0.47  0.47 

but  the  two  expressions  are  so  nearly  identical  that  we  see 
at  a  glance  that  the  ionic  equilibrium  of  the  salt  will  not  be 


INTRODUCTION  27 

affected  appreciably  by  the  presence  of  the  acid.     In  the  case 
of  the  acetic  acid,  however,  we  have  the  initial  relationship, 


0.004X0x34  t  •          j    £  0.004X0.004     ,     0. 

-  ^        °^  =  133  k.  instead  of  -  -  =  k.    Since,  at  equi- 

0.996  -0.996 

librium,  the  fraction  -^  -  -2^-   remains  constant,  and  since, 


owing  to  the  low  H+-ion  concentration,  CHC^O,  cannot  be  in- 
creased appreciably,  nor  Cc2H3oa-  be  appreciably  diminished, 
by  the  formation  of  the  molecular  acid,  it  follows  that  the  value 
of  CH+  must  be  lowered  to  about  y^  of  its  initial  magnitude. 
That  is  to  say,  the  sodium  acetate  in  this  solution  diminishes 
the  hydrogen-ion  concentration  from  0.004  to  about  0.00003. 

The  student  should  especially  note  that  the  concentration  of 
a  given  ion  can  be  lowered  in  this  way  to  a  value  approximating 
zero  only  when  that  ion  unites  with  an  ion  added  to  form  a  substance 
which  is  insoluble  or  which  by  nature  has  only  a  very  slight  tendency 
to  dissociate.  We  might  add  sodium  chloride  in  the  hope 
of  repressing  the  ionization  of  hydrochloric  acid  ;  but,  since 
both  compounds  ionize  highly,  we  should  obtain  no  appreciable 
effect.  If,  however,  we  add  sodium  acetate  in  excess  to  hydro- 
chloric acid,  we  can  obtain  a  solution  which  is  as  weakly  acid 
as  the  one  discussed  above.1 

The  Solubility  Product.  One  of  the  commonest  and  most 
interesting  applications  of  the  law  of  mass  action  is  met  with 
in  connection  with  the  precipitation  and  solution  of  relatively 
insoluble  salts. 

Every  substance  possesses,  when  immersed  in  a  liquid,  a 
certain  solution-tension,  by  which  is  meant  an  expansive  force 
which  tends  to  drive  particles  of  the  substance  outward  into 
the  liquid.  These  particles  move  in  every  direction,  and  conse- 
quently some  of  them  return  to  the  solid  and  reattach  them- 

1  For  example,  i  mol  of  HCl-f  2  mols  of  NaC2H3O2,  in  a  volume  of  i  liter,  give 
a  solution  which  is  uni-molal  in  respect  to  acetic  acid,  to  sodium  acetate,  and  to 
sodium  chloride.  The  hydrogen-ion  concentration  of  this  solution  would  also 
approximate  0.00003. 


28     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

selves  to  it.  This  occurs  the  more  and  more  frequently,  as  the 
concentration  of  the  particles  in  the  liquid  increases,  until, 
finally,  a  stage  is  reached  at  which  the  number  of  particles 
leaving  the  solid  per  unit  of  time  is  equal  to  the  number  deposited 
upon  it.  When  the  entire  liquid  is  equally  charged  with  dissolved 
particles,  the  liquid  immediately  surrounding  the  solid  will  lose 
none  by  diffusion,  and  a  condition  of  equilibrium  will  be  estab- 
lished. At  any  given  temperature,  the  quantity  of  dissolved 
solute  will  remain  thereafter  unchanged,  no  matter  how  long  the 
materials  are  left  in  contact.  It  is  at  this  point  that  the  solution 
is  said  to  be  saturated  with  the  dissolved  substance. 

In  the  case  of  silver  bromate  in  water,  we  have  the  following 
scheme  of  equilibria  : 

AgBrO3  ^±1  AgBr03  ^±:  Ag++Br03- 

(solid)        (dissolved) 

The  solid  AgBr03  molecules  tend  to  enter  the  solution,  while  at 
the  same  time  dissolved  AgBr03  molecules  tend  to  come  out  of 
solution,  and  the  solution  is  saturated  when  these  tendencies 
produce  equal  effects.  The  ions  themselves  (and  any  foreign 
materials  present)  are  not  supposed  to  take  any  direct  part  in 
the  equilibrium  which  controls  solubility.  That  is,  in  solutions 
saturated  at  a  given  temperature  by  a  given  solute,  the  concentra- 
tion of  the  non-ionized  molecules  will  be  constant  no  matter  what 
other  substances  may  be  present,  provided  that  the  quantities  of  all  the 
dissolved  substances  are  not  sufficient  to  alter  the  nature  of  the  solvent. 
The  total  solubility  of  an  ionogen,  as  we  ordinarily  use  the 
term,  is  made  up  of  a  molecular  and  an  ionic  part.  The  quantity 
of  the  latter  does  not  remain  constant  when  a  foreign  substance  giv- 
ing a  common  ion  is  added  to  the  solution.  In  a  solution  of  silver 
bromate,  for  example,  we  have  the  mathematical  relationship  : 


=k'  °r  (Ag+)  x(Br03-)  =*  x(AgBrOs). 

But,  since,  in  the  special  case  of  a  solution  which  is  saturated  with 
the  salt  at  a  given  temperature,  the  concentration  of  the  non- 


INTRODUCTION 


29 


ionized  molecules,  (AgBrOs),  remains  constant,  it  follows  that 
the  product,  &x(AgBr03),  also  remains  constant,  or  that  in  a 
saturated  solution  of  a  given  slightly  soluble  ionogen  the  product 
of  the  concentrations  of  its  ions  is  constant.  This  product  is  called 
the  solubility  product,  because  the  two  separate  values  jointly 
determine  the  magnitude  of  the  total  solubility  of  the  ionogen. 
The  concentration  of  the  non-ionized  molecules  cannot  be 
diminished,  but  the  ionic  part  of  the  solute  may  become  vanish- 
ingly  small  if  the  concentration  of  the  common  ion  is  made  great 
as  compared  with  that  of  the  other  ion  of  the  solute.  The 
relationships  which  exist  in  the  case  of  silver  bromate  are  illus- 
trated in  the  following  table,  where  it  will  be  seen  that  the  ex- 
perimental values  agree  remarkably  well  with  the  calculated  ones. 

SOLUBILITY  OF  AcBROs  IN  MOLS  PER  LITER 


MOLS  PER  LITER  OF 
COMMON-ION  SALT 
ADDED 

SOLUBILITY  FOUND 

SOLUBILITY  CALC. 

Addition  of  Silver 
Nitrate 

Addition  of  Potassium 
Bromate 

(Addition  of  either 
Salt) 

O 
0.00850 
0.0346 

O.OoSlO 
O.OO5IO 
O.O02I6 

0.008  1  o 

O.OO5I9 
0.00227 

0.00504 
0.00206 

The  theory  of  the  precipitation  and  solution  of  slightly  soluble 
ionogens  may  be  summed  up  as  follows : 1 

1  That  is,  of  uni-univalent  ionogens.  In  other  cases,  the  solubility  product 
would  often  contain  ion-concentrations  raised  to  the  second,  third,  etc.,  powers; 
but  in  reality  the  question  is  very  much  complicated  by  interfering  reactions. 
Thus,  in  the  case  of  PbCU,  if  Nad  is  added  to  the  saturated  solution,  some  PbCU 
will  be  precipitated  in  accordance  with  the  theory ;  but  the  addition  of  Pb(NOs)2 
actually  increases  the  solubility  of  the  PbCU-  This  is  probably  because  of  the 
formation  of  complexes,  or  of  intermediate  ions,  such  as  PbCl~,  or  of  both,  whereby 
the  addition  of  the  salt  giving  the  common  bivalent  ion  may  not  only  fail  to  increase 
the  concentration  of  the  bivalent  ion,  but  may  even  lower  that  of  the  univalent  ion. 
At  any  rate,  enough  is  known  to  indicate  that  the  theory  may  not  be  so  much  at 
fault  as  we  ourselves,  in  our  lack  of  methods  for  finding  out  just  what  ions  and 
complexes  are  present  in  such  solutions,  and  in  what  quantities. 


30     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

If  the  product  of  the  concentrations  of  any  pair  of  ions  in  a 
solution  is  made  to  exceed  the  solubility  product  of  the  ionogen 
formed  by  their  union,  the  latter  will  be  precipitated  until  the  ion- 
concentration  product  has  been  reduced  to  its  solubility-product 
value.  And  conversely,  if  the  ion-concentration  product  of  any 
pair  of  ions  in  a  solution  is  made  less  than  the  solubility  product 
of  the  ionogen  formed  by  their  union,  the  latter  will,  if  present  in 
sufficient  excess,  continue  to  dissolve  until  the  ion-concentration 
product  has  been  increased  to  its  solubility-product  value. 

III.   FILTRATION  AND  THE  WASHING  OF  PRECIPITATES 

The  purpose  of  nitration  is  to  separate  a  solid  from  a  liquid 
in  which  it  is  suspended.  This  is  effected  by  causing  the  liquid 
to  pass  through  a  porous  medium  compact  enough  to  retain  the 
solid.  The  most  important  media  in  use  are  filter  paper,  as- 
bestos pulp,  and  platinum  sponge. 

The  Selection  and  Use  of  Paper  Filters.  Three  qualities 
which  are  desirable  in  a  filter  are :  (i)  porosity,  to  insure  rapid 
nitration;  (2)  sufficient  compactness,  to  insure  complete  re- 
tention of  the  precipitate;  and  (3)  low  amount  of  ash.  In 
quantitative  work,  only  filters  should  be  used  which  have  been 
treated  with  hydrochloric  and  hydrofluoric  acids,  and  which, 
on  incineration,  leave  a  small  and  definitely  known  weight  of 
ash.  Such  filters  are  readily  obtainable  in  the  market. 

Rapid  (porous)  filters  should  be  used  for  all  precipitates  which 
do  not  readily  pass  through  the  paper ;  the  slow,  compact  papers 
should  be  used  only  when  necessary.  A  tremendous  amount  of 
time  is  consumed,  often  wasted,  in  the  filtration  and  washing  of 
precipitates. 

The  size  of  the  filter  paper  should  be  determined  by  the  magni- 
tude of  the  precipitate,  and  not  by  the  volume  of  the  liquid  to 
be  filtered.  A  precipitate  should  not  fill  the  paper  more  than 
half  full,  but  if  too  large  a  paper  is  used,  time  is  wasted  in 
washing  the  filter.  The  filter,  as  well  as  the  precipitate,  has  the 
property  of  retaining  certain  salts  very  tenaciously. 


INTRODUCTION  31 

Funnels  should  be  selected  which  have  an  angle  of  60°,  with  a 
narrow  stem  about  eight  niches  long.  The  filter  should  be 
accurately  folded  to  fit  the  funnel,  and  the  top  of  the  filter  should 
be  at  least  one  half  centimeter  below  the  edge  of  the  funnel. 
On  no  account  should  the  paper  project  beyond  the  edge  of  the 
funnel. 

Place  the  filter  in  the  funnel,  wet  it,  and  carefully  bed  it  against 
the  walls  of  the  funnel.  When  the  filter  is  filled  with  distilled 
water,  the  stem  of  the  funnel  should  fill  with  a  column  of  water,1 
and  air  should  not  pass  between  the  funnel  and  the  paper  as 
the  latter  empties.  When  the  filter  is  properly  bedded,  water 
should  flow  through  it  quickly,  and  filtration  will  usually  pro- 
ceed quite  rapidly ;  at  any  rate,  the  paper  will  then  do  its  best. 
The  liquid  at  the  apex  of  the  filter  is  under  a  pressure  approxi- 
mately equal  to  the  weight  of  a  column  of  water  of  the  same 
diameter  as  the  bore  of  the  stem  and  of  a  height  equal  to  the 
length  of  the  stem  plus  the  depth  of  liquid  in  the  filter. 

When  paper  filters  are  employed,  the  use  of  a  vacuum  pump 
to  promote  filtration  is  of  doubtful  advantage  in  quantitative 
analysis.  The  increased  tendency  of  precipitates  to  pass  through 
the  filter  more  than  offsets  the  possible  gain  in  time.  Whenever 
suction  is  applied,  a  more  compact  paper  should  be  used,  or  the 
point  of  the  filter  should  be  supported  by  a  perforated  platinum 
cone. 

The  vessel  used  to  receive  the  initial  filtrate  should  invariably 
be  replaced  by  a  clean  one,  properly  labeled,  before  the  washing 
of  a  precipitate  is  begun.  Precipitates  which  at  first  show  no 
tendency  to  pass  through  the  filter  may  enter  into  colloidal  solu- 
tion as  the  washing  proceeds.  The  advantage  gained  in  such  an 
instance  by  having  removed  the  first  filtrate  is  obvious. 

The  precipitate  is  generally  allowed  to  settle  before  filtration. 
The  clear  liquid  should  not  be  poured  directly  on  to  the  filter, 

1  If  this  fails  to  take  place,  either  the  stem  of  the  funnel  is  too  wide,  or  it  is  not 
free  from  grease.  The  latter  can  be  removed  by  means  of  warm  sulphuric  acid- 
dichromate  cleaning  solution. 


32     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

but  down  a  glass  rod,  from  which  the  stream  should  be  directed 
towards  the  side,  not  the  center  of  the  filter  paper.  The  re- 
ceiving vessel  for  the  nitrate  should  be  placed  so  that  the  liquid 
from  the  funnel  will  run  down  its  side ;  otherwise  there  is  danger 
of  loss  by  splashing.  As  far  as  possible,  the  precipitate  itself 
should  be  kept  in  the  precipitation  vessel. 

Much  time  can  generally  be  saved  by  washing  the  precipitate 
by  decantation,  i.e.  by  adding  to  it  successive  portions  of  wash 
liquid,  allowing  the  precipitate  each  time  to  settle,  and  decant- 
ing the  clear  supernatant  liquid  through  the  filter,  without  un- 
duly disturbing  the  precipitate.  This  procedure  is  especially 
advantageous  in  the  case  of  precipitates  which  tend  to  clog  the 
pores  of  the  filter.  Finally  the  precipitate  may  be  transferred 
to  the  filter,  and  the  washing  completed  there. 

It  will  always  be  found  that  small  portions  of  the  precipitate 
adhere  to  the  walls  and  bottom  of  the  containing  vessel.  These 
can  be  rubbed  loose  by  means  of  a  so-called  "  policeman,"  —  a 
piece  of  soft  rubber  tubing  tightly  fitted  on  the  end  of  a  glass 
rod.  Pieces  of  rubber  tubing  with  closed  ends  are  sold  for  the 
purpose.  These  rubber-tipped  rods  should  be  used  only  for  the 
above-mentioned  purpose;  they  should  never  be  allowed  to  stand 
in  analytical  solutions,  nor  should  they  be  used  as  ordinary 
stirring  rods. 

Precipitates  should  never  be  allowed  to  dry  before  they  have 
been  completely  washed.  They  are  likely  to  shrink  and  crack, 
and,  on  further  washing,  the  liquid  will  pass  through  these 
channels  only ;  this  is  especially  true  of  gelatinous  precipitates. 

Every  original  filtrate  must  be  properly  tested  to  insure  com- 
plete precipitation,  and  the  wash  waters  also  must  be  examined. 
It  is  useless,  however,  to  test  the  latter  until  several  washings 
have  been  made.  Only  a  few  drops  should  be  taken  if  the  filtrate 
is  to  be  used  for  a  subsequent  determination;  but  when  the 
washing  is  nearly  finished,  at  least  2  or  3  cc.  should  be  used. 
The  necessity  of  making  these  tests  cannot  be  too  strongly  impressed 
upon  the  student;  and  no  exception  should  ever  be  made. 


INTRODUCTION  33 

Wash  Bottles.  Wash  bottles,  for  distilled  water,  should  con- 
sist of  flasks  of  about  500  cc.  capacity  which  are  provided  with 
rubber  stoppers  and  with  tubes  gracefully  bent  and  not  too  long. 
The  jet  should  be  connected  in  a  flexible  manner  with  the  outlet 
tube  by  means  of  a  short  piece  of  soft  rubber  tubing,  and  should 
deliver  a  smooth  stream  about  i  mm.  in  diameter.  For  use  with 
hot  water,  the  neck  of  the  bottle  should  be  wrapped  with  heavy 
asbestos  twine,  or  other  suitable  material.  In  order  to  avoid 
mistakes,  wash  bottles  for  other  liquids  than  distilled  water 
should  always  be  plainly  labelled. 

Gooch's  Filtration  Crucible.  In  1878,  F.  A.  Gooch  proposed 
the  separation  of  certain  precipitates  by  nitration  with  suction 
through  a  mat  of  asbestos  bedded  on  the  perforated  bottom  of  a 
crucible.  The  precipitate  then  could  be  washed,  dried,  and 
weighed  in  the  crucible. 

Preparation  of  the  Asbestos.  There  are  several  varieties  of 
asbestos  in  the  market,  of  which  the  long-fiber  "  silky  "  chrys- 
olite asbestos  is  the  best.  Rub  the  asbestos  roughly  over  the 
surface  of  a  jo-mesh  brass  sieve,  placed  in  an  inverted  position 
on  a  sheet  of  paper,  until  a  sufficient  quantity  has  passed  through. 
Shake  this  up  with  water,  allow  most  of  it  to  settle,  and  pour 
off  the  very  fine  particles.  Now,  digest  the  pulp  on  the  steam 
bath  for  i  hour  with  strong  hydrochloric  acid,  in  a  covered 
porcelain  dish.  At  the  end  of  this  operation,  dilute  the  mixture 
with  water,  pour  off  the  liquid  through  a  funnel  provided  with 
a  platinum  filtration  cone,  and  wash  the  asbestos  with  hot  water, 
at  first  by  decantation,  and  finally  in  the  funnel,  using  gentle 
suction,  until  a  5  cc.  portion  of  the  filtrate  fails  to  give  an  opales- 
cence  with  silver  nitrate.  Mix  the  washed  asbestos  with  dis- 
tilled water  and  keep  it  in  a  bottle  ready  for  use. 

Packing  the  Crucible  with  Asbestos  Felt.  Stretch  a  piece  of 
rubber-band  tubing  over  the  upper  edges  of  a  cylindrical  glass 
tube,  about  3  cm.  in  diameter  and  7  cm.  long,  which  is  closed  at 
one  end  except  for  an  attached  stem  of  suitable  size  and  length 
to  pass  through  a  rubber  stopper.  The  tubing  should  project 


34     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

1.5-2  cm.  above  the  top  of  the  funnel  tube.  Fit  the  glass  funnel 
tube  into  the  stopper  of  a  filter  bottle,  connect  the  latter  with 
the  vacuum  pump,  and  then  press  into  the  short  projecting  end 
of  the  rubber  tube  the  Gooch  Crucible,  so  that  it  fits  in  an  air- 
tight manner. 

Take  some  of  the  asbestos  suspension  referred  to  above,  add 
water,  and  stir  the  mixture.  Allow  this  to  settle  for  some  time, 
pour  off  the  very  fine  particles,  apply  a  gentle  suction,  and  then 
pour  some  of  the  mixture  cautiously  into  the  crucible  until  an 
even  felt  of  asbestos,  not  over  1.5  mm.  in  thickness,  is  formed. 
Place  a  small  perforated  disk  (filter  plate)  upon  the  asbestos,  and 
pour  just  enough  more  asbestos  into  the  crucible  to  barely  cover 
the  disk.  Run  water  through  the  crucible  until  no  more  asbestos 
fibers  run  through,  and  make  sure  that  the  washings  are  free 
from  chlorides.  If  the  water  which  has  passed  through  the  cruci- 
ble is  held  before  a  bright  light,  any  suspended  asbestos  fibers 
can  readily  be  seen.  Usually,  250-500  cc.  of  water  will  suffice. 
The  perforated  filter  plate  is  used  to  protect  the  asbestos  felt, 
during  the  washing  and  subsequent  filtration. 

Place  the  crucible  in  a  small  beaker,  dry  at  120-130°  for  an 
hour,  cool  in  a  desiccator,  and  weigh.  Heat  again,  for  ^  hour, 
cool,  and  again  weigh,  repeating  until  the  weight  is  constant 
within  0.0002  g.  The  filter  is  then  ready  for  use. 

How  to  Use  the  Gooch  Crucible.  The  weighed  crucible  is  re- 
placed in  the  funnel,  and  a  gentle  suction  applied,  after  which 
the  liquid  to  be  filtered  may  be  passed  through  the  crucible,  and 
the  precipitate  washed  as  if  the  crucible  were  a  filter  paper  and 
funnel.  When  pouring  liquid  into  the  crucible,  hold  the  stirring 
rod  well  down  in  the  crucible,  so  as  not  to  disturb  the  asbestos. 
Always  examine  the  first  portions  of  the  filtrate  with  great  care 
for  asbestos  fibers,  and  refilter  the  liquid  if  any  are  visible. 

When  the  precipitate  has  been  washed  and  the  crucible  dried, 
the  whole  is  weighed.  The  drying  and  weighing  should  be  re- 
peated, as  above,  to  constant  weight.  The  increase  in  weight 
represents  the  weight  of  the  precipitate. 


INTRODUCTION  35 

The  same  crucible  can  be  used  for  a  number  of  determinations 
of  the  same  substance.  When  the  collection  of  precipitates  in 
the  crucible  becomes  too  large,  the  upper  part  can  be  removed, 
and  the  crucible  used  as  before.  If.  the  felt  has  been  properly 
prepared,  filtration  and  washing  are  rapidly  accomplished,  and 
this,  combined  with  the  possibility  of  repeatedly  using  the  same 
filter,  is  a  strong  argument  in  its  favor,  with  any  but  gelatinous 
precipitates. 

If  perforated  platinum  crucibles  are  available,  which  after 
removal  from  the  funnel,  can  be  fitted  into  platinum  cups,  the 
precipitates  can  be  ignited  as  in  ordinary  crucibles.  In  this  case, 
however,  it  is  better  to  pack  the  Gooch  crucible  with  a  felt  of 
platinum  sponge.  In  this  form,  the  apparatus  is  known  as  a 
Munroe  Crucible. 

Sources  of  Error.  It  is  important  to  remember  that  asbestos 
may  absorb  appreciable  amounts  of  alkali,  not  removed  by 
washing,  so  that,  as  a  rule,  solutions  containing  fixed  alkalies 
should  not  be  filtered  through  asbestos.  Asbestos  is  also  slightly 
attacked  by  water  and  feebly  acid  solutions ;  but  after  the  treat- 
ment indicated  above,  there  is  no  real  danger  from  this  source. 
If  the  felt  has  been  properly  prepared,  there  is  no  danger  of  losing 
asbestos  during  the  filtration  and  washing,  but  the  liquid  which 
runs  through  should  nevertheless  be  carefully  examined,  and 
refiltered  if  necessary. 

The  Theory  of  Washing  Precipitates.  The  theory  of  wash- 
ing precipitates  should  include  the  consideration  of  several 
factors,  among  which  may  be  mentioned  the  phenomena  of 
adsorption  (in  which  the  filter  also  takes  part),  and  the  tend- 
ency of  precipitates  to  enter  the  wash  liquid  in  colloidal  form. 
But  these  factors  have  been  discussed  under  precipitation. 
Aside  from  these  considerations,  there  is  the  important  ques- 
tion concerning  the  most  effective  method  of  washing  precipitates. 

Let  us  suppose  that  a  precipitate  is  to  be  washed  by  decanta- 
tion,  and  for  the  sake  of  simplicity  let  us  assume  that  neither 
it  nor  the  filter  exercises  any  physical  or  chemical  action  on  the 


36     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

salts  dissolved  in  the  mother  liquid.  The  supernatant  mother 
liquid  has  been  decanted  as  far  as  possible  into  the  filter,  and 
the  latter  has  been  allowed  to  drain. 

Let  v  cc.  be  the  total  volume  of  solution  which  remains  in 
contact  with  the  precipitate  and  filter,  and  V  cc.  the  volume  of 
wash  water  added  each  tune  ;  and  assume  that  the  latter  mixes 
uniformly  with  the  liquid  adhering  to  the  precipitate  and  filter. 
Then,  upon  the  addition  of  V  cc.  of  water,  the  total  volume  of 
liquid  is  (V+v)  cc.  Further,  let  Co  be  the  concentration  in 
grams  per  cubic  centimeter  of  the  undesirable  salts  in  the 
original  solution;  then  the  quantity  contained  in  the  v  cc.  left 
in  contact  with  the  precipitate  and  filter  is  vC0  grams.  By 
the  addition  of  V  cc.  of  water,  the  concentration  is  reduced  to 

Ci=  -r-Co,  and  if  this  liquid  is  removed  until  only  v  cc.  are 


left,  the  quantity  of  undesirable  salts  present  is  reduced  to 
flCo.    A  second  addition  of  V  cc.  of  water  gives  the 


ntrati 
salts  in  the  v  cc.  left  on  draining  is  now 


concentration,  C2  =  ~—  Ci  =  f  —  ~-  J  C0,  and   the  quantity   of 


or,  after  n  washings,  the  quantity  of  undesirable  salts  has  been 
diminished  to  the  value, 


This  formula  expresses  mathematically  the  self-evident  fact 
that,  for  a  given  number  of  washings,  the  quantity  of  undesir- 
able salts  left  behind  will  be  the  smaller,  the  more  completely 
the  precipitate  and  filter  are  drained,  and  the  greater  the  volume 
of  the  wash  water  that  is  added  each  time.  The  formula  enables 
us,  however,  to  find  the  answer  to  a  less  simple  question;  viz., 
What  is  the  most  efficient  method  of  washing  a  precipitate  with 


INTRODUCTION  37 

a  given  amount  of  wash  liquid  ?  Suppose,  for  example,  we  wish 
to  use  only  150  cc.  of  wash  liquid  :  is  it  better  to  wash  six  times 
with  25  cc.  portions,  or  to  wash  10  times  with  15  cc.  portions? 
Let  us  assume  that  C0  =  o.i  g.  per 'cubic  centimeter,  and  that 
v  =  5  cc. ;  then,  in  the  two  cases,  the  quantities  of  undesirable 
salts  left  behind  will  be  ( A) 6X 0.5  =  0.0000 107  g.,  and  (A)10Xo.5 
=  0.00000047  g.,  respectively.  Disregarding  adsorption,  which 
greatly  decreases  the  efficiency  of  washing,  ten  washings  with 
15  cc.  portions  are  23  times  as  efficient  as  six  washings  with 
25  cc.  portions.  Both  methods  of  procedure  will  require  ap- 
proximately equal  intervals  of  time,  since,  in  either  case,  153  cc. 
of  liquid  must  run  through  the  filter.  //  is  much  better  to  wash 
a  precipitate  many  times  with  small  portions  of  liquid,  than  a  few 
times  with  larger  portions.  Each  portion  of  wash  liquid  should 
be  removed  as  far  as  possible  by  decantation  and  drainage,  before 
the  addition  of  a  fresh  portion. 

Another  factor  to  be  considered  is  the  temperature  of  the 
solution  to  be  filtered.  Since  the  rate  of  flow  of  a  liquid  through 
a  filter  depends  largely  upon  the  viscosity  of  the  liquid,  and 
since  the  viscosity  of  water  at  100°  is  only  one  sixth  that  of  water 
at  o°,  it  is  well  to  filter  and  wash  at  a  high  temperature,  unless 
there  is  good  reason  to  the  contrary. 

Finally,  in  washing  a  precipitate  on  a  paper  filter,  great  care 
must  be  taken  to  wash  the  filter  itself.  Soluble  salts  are  tena- 
ciously held  back  at  the  upper  edges  of  the  paper,  and  therefore 
this  part  of  the  filter  should  receive  especial  attention.  It  is 
best  to  fill  the  filter  each  time,  and,  before  refilling,  to  allow  it  to 
drain  completely.  This  is  a  strong  argument  in  favor  of  filters 
which  are  not  too  large. 

IV.  THE  DRYING  AND  IGNITION  OF  PRECIPITATES 

Drying  Ovens.  There  are  on  the  market  many  types  of 
drying  ovens,  heated  by  gas,  by  steam  pipes,  or  by  elec- 
tricity, in  which  the  temperature  may  be  more  or  less  accu- 


38     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

rately  controlled.  The  oven  consists  essentially  of  a  drying 
chamber,  through  which  there  is  provided  a  slow  circulation 
of  hot  air. 

A  precipitate  is  dried  on  the  filter  by  placing  the  funnel  con- 
taining both  in  a  drying  oven,  at  90-100°,  and  leaving  it  there 
for  a  sufficient  time.  The  funnel  should  be  covered  with  a 
sheet  of  common  filter  paper,  fastened  in  place  by  crimping  its 
edges  over  those  of  the  funnel.  If  the  precipitate  is  suitable 
for  weighing  without  ignition,  e.g.  silver  chloride  in  a  Gooch 
crucible,  it  should  be  dried  to  constant  weight  at  a  temperature 
considerably  above  the  boiling  point  of  water  (in  this  case,  at 
120-130°),  in  order  to  remove  the  last  traces  of  moisture  from 
the  filter  as  well  as  from  the  precipitate.  (Before  it  is  used,  the 
packed  Gooch  crucible  should,  of  course,  be  dried  to  constant 
weight  at  the  same  temperature.) 

Many  precipitates  may,  under  proper  precautions,  be  ignited 
without  previous  drying  in  an  oven.  But  if  such  precipitates 
can  be  dried  over  night,  or  otherwise,  without  loss  of  time  to 
the  analyst,  it  is  well  to  submit  them  to  this  process. 

The  precipitate,  with  the  filter  folded  over  it,  is  placed  at  the 
bottom  of  the  crucible,  and  the  latter  is  supported,  on  a  tri- 
angle, so  far  above  the  small  flame  of  a  burner  as  to  preclude 
the  violent  escape  of  steam.  When  the  filter  and  contents  are 
dry,  the  open  crucible  is  tilted  on  its  side,  and  the  heat  slightly 
increased  until  the  filter  chars;  the  heating  is  then  continued 
at  this  rate  until  the  gases  from  the  dry  distillation  of  the  paper 
have  been  completely  expelled  without  taking  fire.  In  this 
way,  no  material  will  be  lost  owing  to  strong  draughts  caused 
within  the  crucible  by  burning  gases. 

During  the  dry  distillation  of  the  paper,  the  flame  should  be 
placed  near  the  mouth  of  the  crucible ;  but  afterwards  it  should 
be  well  at  the  base  of  the  inclined  crucible,  to  allow  a  ready  access 
of  air.  After  the  filter  has  been  freed  from  volatile  matter,  the 
crucible  should  be  heated  to  redness  until  the  ignition  is  com- 
plete. 


UNIVERSITY  OF  CALIFORNIA 
DEPARTMENT   OF   CIVIL   ENGINEER1N 


Some  precipitates  are  reduced  or  otherwise  affected  by  con- 
tact with  hot  carbon  or  reducing  gases  from  the  filter  paper; 
e.g.  silver  chloride,  lead  sulphate,  etc.  are  reduced  to  metal. 
Since,  however,  these  metals  are  volatile  only  at  very  high  tem- 
peratures, there  is  no  loss  in  their  case,  and  the  metal  can  by 
suitable  treatment  be  transformed  quantitatively  into  the 
original  compound.  In  such  cases  it  is  advisable  to  separate 
the  precipitate  as  far  as  possible  from  the  filter,  and  then  to  ignite 
the  latter.  The  small  quantity  of  reduced  metal  is  moistened 
with  a  few  drops  of  nitric  acid,  and  the  resulting  nitrate  con- 
verted into  silver  chloride  with  hydrochloric  acid,  or  into  lead 
sulphate  with  sulphuric  acid,  and  the  excess  of  acid  expelled  by 
cautiously  heating  the  crucible.  The  bulk  of  the  precipitate  is 
then  added,  and  the  whole  ignited. 

Unless  specially  directed,  precipitates  should  not  be  heated 
over  the  blast  lamp. 

Desiccators.  After  an  object  has  been  dried  and  ignited,  it 
must  be  permitted  to  cool  before  it  can  be  weighed  with  accu- 
racy. In  order  to  protect  it  from  contamination  with  moisture, 
carbon  dioxide,  etc.,  it  should  invariably  be  allowed  to  cool  in  a 
desiccator. 

For  general  analytical  work,  desiccators  should  be  charged 
with  fragments  of  fused,  anhydrous  calcium  chloride,  some 
distance  above  which  is  placed  a  porcelain  plate  provided  with 
holes  of  a  size  suitable  for  the  reception  of  crucibles.  In  order 
to  give  the  cover  of  the  desiccator  an  air-tight  fit,  the  ground- 
glass  contact  surfaces  should  be  thinly  coated  with  vaseline  or 
some  similar  substance.1 

Desiccators  should  never  be  left  uncovered.  The  dehydrat- 
ing agent  is  intended  to  keep  the  air  inside  the  desiccator  dry  ; 
it  rapidly  loses  its  efficiency  if  exposed  to  the  outside  air.  If  the 
lumps  of  calcium  chloride  tend  to  stick  together,  the  charge 
should  be  renewed. 

1  A  mixture  made  by  melting  together  equal  parts  of  vaseline  and  beeswax  is 
very  suitable. 


40     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Pumice  moistened  with  concentrated  sulphuric  acid  is  some- 
times used  instead  of  calcium  chloride. 

Crucibles.  The  most  commonly  used  crucibles  are  of  high 
grade  porcelain.  They  withstand  very  high  temperatures 
without  appreciable  change  of  weight,  and  they  are  compara- 
tively cheap  in  price.  They  cannot  be  used  for  fusions  because 
most  fluxes,  particularly  those  of  a  basic  nature,  attack  the 
glaze  as  well  as  the  porcelain  itself.  Even  in  the  ignition  of 
non-basic  precipitates,  in  spite  of  the  most  careful  washing, 
traces  of  fusible  materials  always  remain  with  the  precipitate, 
and  these  in  time  destroy  the  glaze.  After  the  lining  is  thus 
roughened  it  is  difficult  to  clean,  and  the  crucible  is  unsuitable 
for  use. 

Crucibles  more  or  less  suitable  for  the  ignition  of  precipitates 
are  also  made  of  alundum,  and  of  fused  silica. 

Crucibles  of  platinum  are  very  desirable  for  ignitions;  and 
for  many  fusions  they  are  essential.  Platinum  melts  at  about 
1770°  and  does  not  soften  enough  to  preclude  its  use  at  tem- 
peratures slightly  below  its  melting  point.  It  is  soluble  in  liquids 
containing  free  chlorine,  such  as  nitrate-chloride  mixtures  of 
acid  reaction,  and  to  a  lesser  degree  in  acid  ferric  chloride  solu- 
tions. Care  must  be  taken  to  prevent  injury  to  platinum  vessels, 
or  the  introduction  of  platinum  into  solutions,  by  a  disregard  of 
these  facts. 

Platinum  easily  alloys  with  most  metals,  and  for  that  reason  it 
should  not  ordinarily  be  heated  in  contact  with  metals,  or  with 
compounds  of  easily  reducible  metals,  —  never,  if  carbon  or 
reducing  gases  are  also  present.  When  heated  for  a  long  time 
in  contact  with  carbon,  platinum  slowly  takes  up  the  latter  and 
becomes  brittle;  therefore,  the  crucible  should  never  be  heated 
in  a  reducing  flame.  The  flame  should  be  carefully  adjusted  so 
that  the  tip  of  the  inner  cone  is  below  the  bottom  of  the  crucible, 
and  a  flame  showing  yellow  must  never  be  used.  Compounds 
of  phosphorus  or  arsenic  must  not  be  heated  under  reducing 
conditions,  since  the  free  elements,  as  well  as  phosphides  and 


INTRODUCTION  41 

arsenides,  render  platinum  brittle  and  lower  its  melting  point. 
"  Unknown  "  substances  should  never  be  heated  in  platinum 
vessels. 

Platinum  ware  should  always  be  kept  bright  and  clean.  For 
this  purpose  it  should  be  frequently  polished  with  fine  sea  sand 
or  with  precipitated  silica.  These  remove  most  impurities,  and 
polish  the  platinum  without  serious  loss.  The  fusion  of  potas- 
sium bisulphate  in  the  vessel  is  a  good  method  for  cleaning  the 
badly  tarnished  inside.  The  bisulphate  should  be  poured  out 
of  the  crucible  while  still  liquid ;  for  if  it  has  been  strongly  heated, 
the  melt  (pyrosulphate)  is  apt  to  expand  so  rapidly  on  cooling 
as  to  burst  the  crucible. 

Never  heat  the  platinum  crucible  or  dish  in  contact  with  iron 
or  metals  other  than  platinum,  nor  place  hot  platinum  in  con- 
tact with  foreign  metals.  Use  nothing  but  pipeclay,  quartz, 
or  platinum  triangles,  and  platinum-shod  tongs.1 

V.   THE  EVAPORATION  OF  LIQUIDS 

The  greatest  care  must  be  taken  to  prevent  loss  of  material 
during  processes  of  solution  and  evaporation,  either  from  the 
evolution  of  gas,  from  too  violent  ebullition,  or  from  evaporation 

1  Modern  platinum  ware  is  often  inferior  in  quality  to  that  on  the  market  some 
years  ago,  and  the  cause  has  been  the  subject  of  special  inquiry  by  a  committee  of 
the  American  Chemical  Society.  The  main  objections  are:  "(*)  Undue  loss  of 
weight  on  ignition ;  (2)  undue  loss  on  acid  treatment,  especially  after  strong  igni- 
tion ;  (3)  unsightly  appearance  of  the  surface  after  strong  ignition,  especially  after 
the  initial  stages  of  heating;  (4)  adhesion  of  crucibles  and  dishes  to  triangles, 
sometimes  to  such  an  extent  as  to  leave  indentations  on  the  vessel  at  the  points 
of  contact  with  the  triangle,  even  when  complete  cooling  has  been  reached  before 
the  two  are  separated ;  (5)  alkalinity  of  the  surface  of  the  ware  after  strong  ignition ; 
(6)  blistering;  and  (7)  development  of  cracks  after  continued  heating."  It  is  the 
general  opinion  that  the  trouble  arises  from  the  working  of  scrap  platinum  into 
chemical  ware.  The  main  difficulties  here  mentioned  are  not  characteristic  of 
platinum  ware  from  some  of  the  best  manufacturers. 

The  committee  recommends  that  purchasers  specify  that  platinum  ware  must 
show  no  marked  uneven  discoloration  on  heating,  must  give  no  test  for  iron  after 
heating  for  two  hours,  and  that  the  rate  of  loss  per  hour  at  1100°  over  a  period  of 
four  hours  shall  not  exceed  0.2  mg. 


42     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

nearly  to  dryness  accompanied  by  spattering  or  by  the  "  crawl- 
ing "  of  salts  over  the  edge  of  the  vessel.  In  order  to  prevent 
mechanical  losses,  solutions  in  which  gases  are  being  evolved, 
or  which  are  to  be  boiled  on  the  hot  plate,  should  invariably 
be  covered.  And  liquids  which  contain  suspended  matter 
(precipitates)  should  always  be  cautiously  heated,  since  the 
presence  of  solid  matter  frequently  occasions  violent  "  bumping  " 
which  may  lead  to  mechanical  losses  or  to  the  destruction  of 
the  vessel. 

The  evaporation  of  aqueous  solutions  rarely  requires  the  use 
of  temperatures  above  100°.  Temperatures  somewhat  below 
this  point,  but  sufficient  for  the  evaporation  of  most  aqueous 
solutions,  are  best  attained  by  the  use  of  the  steam  bath,  which 
has  the  advantage  of  keeping  the  solutions  at  a  temperature 
below  that  at  which  mechanical  losses  are  to  be  feared. 

Evaporations  should  not  be  attempted  in  tall,  narrow  vessels, 
but  should  be  carried  out  in  low,  wide-mouthed  dishes  or  casse- 
roles :  it  is  obvious  that  evaporation  is  promoted  by  the  ex- 
posure of  a  large  surface  of  liquid  to  the  air.  In  evaporations 
on  the  steam  bath,  a  watch  glass  should  be  supported  above  the 
casserole  or  dish  by  means  of  a  glass  triangle  or  other  suitable 
device. 

If  a  large  volume  of  liquid  is  to  be  evaporated,  it  is  not  neces- 
sary that  the  vessel  should  contain  it  all  at  once.  Fresh  portions 
of  the  liquid  may  be  added,  from  time  to  tune,  as  the  volume  of 
that  in  the  vessel  is  reduced  by  evaporation. 

Liquids  should  never  be  transferred  from  one  vessel  to  another, 
nor  to  a  filter,  without  the  aid  of  a  stirring  rod  held  firmly  against 
the  lip  or  side  of  the  containing  vessel.  In  order  to  prevent  the 
loss  of  liquid  by  running  down  on  the  outside  of  the  vessel,  a 
very  thin  coating  of  vaseline,  applied  with  the  finger  to  the  out- 
side edge  of  the  vessel,  will  suffice.  If  the  vessel  is  provided 
with  a  lip,  this  is  usually  unnecessary. 

As  few  transfers  of  liquid  as  possible  from  one  vessel  to  another 
should  be  made  during  an  analysis.  In  such  transfers,  the  solu- 


INTRODUCTION  43 

tion  must,  of  course,  be  quantitatively  washed  out.  This  can 
be  accomplished  better  by  the  use  of  successive  small  portions 
of  wash  water,  say  of  5-10  cc.  each,  than  by  the  addition  of  a 
few  larger  portions  which  unnecessarily  increase  the  volume  of 
the  solution  and  lead  to  loss  of  time  in  subsequent  nitrations  or 
evaporations. 

VI.   THE  VOLUMETRIC   MEASUREMENT  OF  LIQUIDS1 

Measurements  with  a  good  balance  and  weights  can  often 
be  made  with  a  precision  even  greater  than  is  necessary  for 
general  analytical  work.  But,  as  has  been  intimated,  the  errors 
involved  in  the  preparation  of  a  troublesome  precipitate  for 
weighing  may  impair  the  value  of  an  exact  weighing.  Although 
the  measurement  of  volume,  in  volumetric  analysis,  is  not  apt 
to  be  so  precise  and  reliable  as  the  measurement  of  weight,2  yet 
the  results  of  volumetric  processes,  based  on  suitable  reactions, 
are  frequently  more  trustworthy  than  those  of  gravimetric 
processes,  because  the  volumetric  process  for  the  determination 
of  the  substance  is  less  liable  to  error.  With  proper  precautions 
many  volumetric  processes  yield  excellent  results ;  and,  es- 
pecially in  technical  work,  where  time  is  an  essential  factor, 
volumetric  processes  are  very  often  used  in  preference  to  gravi- 
metric. In  order,  however,  that  dangerous  errors  may  be- 
eliminated  in  volumetric  work,  it  is  of  great  importance  that 
the  analyst  should  have  a  clear  idea  of  the  precautions  necessary 
for  the  attainment  of  a  high  degree  of  accuracy. 

1  For  more  detailed  information  on  this  subject,  see  Bulletin  of  the  Bureau  of 
Standards,  Vol.  4,  pp.  553-601  (1908). 

2  Even  this  difficulty  can  be  obviated  by  the  use  of  weight  burettes;  i.e.  of  burettes 
of  such  construction  as  to  be  readily  weighable  both  before  and  after  the  removal 
of  the  quantity  of  solution  required  for  the  completion  of  the  given  reaction.     The 
difference  gives  the  weight  of  solution  required,  and,  provided  the  solution  has 
been  standardized  by  the  same  method,  the  quantity  of  the  substance  under  in- 
vestigation can  be  readily  calculated.    In  this  way,  if  based  on  suitable  reactions, 
exceedingly  exact  determinations  can  be  executed.     (A  similar  weight  burette 
should  always  be  used  as  a  counterpoise.) 


44     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Volumetric  Apparatus.  The  exact  measurement  of  liquid 
volumes  necessitates  the  use  of  certain  special  forms  of  apparatus, 
the  most  important  of  which  will  now  be  described. 

Burettes  are  graduated  glass  tubes  of  uniform,  small  diam- 
eter, designed  to  measure  variable  amounts  of  liquids  delivered 
by  them  when  supported  in  a  vertical  position.  The  outflow 
of  the  liquid  is  controlled  either  by  a  glass  stopcock  or  by  means 
of  a  rubber  joint  which  connects  the  end  of  the  tube  with  a  glass 
nozzle,  and  which  is  provided  with  a  pinchcock  or  other  suitable 
device.  The  former  require  the  use  of  some  lubricant,  such  as 
vaseline,  to  permit  easy  control  of  the  outflow,  and  the  latter 
have  the  disadvantage  that  the  rubber  connection  is  acted  upon 
to  some  extent  by  certain  solutions,  which  in  consequence  are 
apt  to  experience  a  change  in  concentration;  e.g.  rubber  stop- 
cocks should  never  be  used  with  permanganate  or  iodine  solu- 
tions. For  accurate  work,  50  cc.  burettes  should  be  graduated 
to  o.i  cc.  and  25  cc.  and  30  cc.  burettes  to  0.05  cc.,  and  the 
graduation  marks  should  be  separated  by  at  least  i  mm. 

Transfer  Pipettes  are  tubes  of  much  smaller  bore  than  bu- 
rettes, designed  to  deliver  specific  volumes  of  liquid.  They  are 
provided  with  an  enlargement  at  the  center,  which  greatly 
reduces  the  length  required,  and  with  a  single  mark  on  the  upper 
length,  which  indicates  the  point  to  which  they  must  be  filled  in 
order  to  deliver  the  indicated  volume  of  liquid.  Pipettes  are 
filled  by  suction  and  are  allowed  to  deliver  the  liquid,  while 
held  in  a  vertical  position,  by  the  action  of  gravity.  On  account 
of  the  smallness  of  the  bore  at  the  point  where  the  reading  is 
made,  the  pipette  may  be  made  to  measure  liquids  with  a  high 
degree  of  accuracy.  However,  certain  errors  of  manipulation 
often  render  the  measurements  inexact. 

Measuring  Flasks  are  employed  for  measuring  relatively 
large  volumes  of  liquid  in  one  portion.  The  neck  must  be 
sufficiently  small  to  permit  a  reading  with  only  a  slight  per- 
centage error,  but  large  enough  to  permit  filling  and  emptying 
without  trouble.  The  neck  should  also  be  of  uniform  bore  and 


INTRODUCTION  45 

of  some  margin  above  and  below  the  mark.  For  the  most 
accurate  work  the  flask  is  always  graduated  for  containing  the 
amount  indicated  by  the  inscription  upon  it. 

Graduated  Cylinders  are  cylindrical  glass  vessels  provided 
with  an  enlarged  base  or  foot  and  with  a  lip  for  pouring.  They 
are  marked  to  indicate  the  varying  amounts  of  liquid  which 
they  may  contain  and  are  employed  for  rough  measurements  only. 

Necessary  Precautions  in  the  Use  of  Volumetric  Apparatus. 
In  making  use  of  volumetric  methods,  the  following  are  the  most 
important  sources  of  error;  they  must  be  fully  reckoned  with 
if  the  results  are  to  be  reliable. 

Errors  Due  to  Water  in  the  Apparatus.  It  is  usually  neces- 
sary to  rinse  the  apparatus  with  water  before  using ;  and  the 
amount  of  water  retained  may  appreciably  change  the  concen- 
tration of  the  solution  measured.  The  error  can  be  avoided  by 
drying  the  apparatus  before  use,  or,  more  conveniently,  by 
rinsing  it  out  with  several  small  portions  of  the  solution  to  be 
measured,  and  discarding  the  washings. 

Errors  Due  to  Drainage  or  Afterflow.  When  a  liquid  is  per- 
mitted to  flow  somewhat  rapidly  from  a  burette,  or  pipette, 
small  amounts  adhere  to  its  inner  surface,  and  gradually  flow 
down  and  unite  with  the  liquid  still  remaining  hi  the  vessel. 
In  order  to  avoid  errors  from  this  source  the  rate  of  outflow 
must  be  limited  by  the  size  of  the  outlet,  and  a  sufficient  interval 
must  elapse  between  the  time  at  which  the  flow  from  the  ap- 
paratus is  stopped,  and  at  which  the  reading  is  made.  This  inter- 
val, unless  indicated  on  the  instrument,  may  be  taken  as  30  seconds 
in  the  case  of  burettes  and  15  seconds  in  the  case  of  pipettes. 

In  the  case  of  transfer  pipettes,  the  outlets  must  be  of  such 
size  that  the  free  outflow  shall  last  not  more  than  one  minute 
and  not  less  than  15,  20,  and  30  seconds  respectively,  for  5,  10, 
and  50  cc.  pipettes.  The  rate  of  outflow  of  burettes  must  not 
be  more  than  three  minutes,  nor  less  than  90  and  50  seconds, 
respectively,  for  50  and  30  cc.  burettes.  Burette  and  pipette 
tips  should  be  made  with  a  gradual  taper  of  2-3  cm. ;  a  sudden 


46     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

contraction  at  the  orifice  is  not  permitted,  and  the  tip  should 
be  well  finished. 

In  filling  pipettes  and  burettes  excess  liquid  adhering  to  the 
tip  should  be  removed  when  completing  the  filling.  In  emptying 
pipettes  and  burettes,  they  should  be  held  in  a  vertical  position, 
and  after  the  continuous  unrestricted  outflow  has  ceased  the 
tip  should  be  touched  with  the  wet  surface  of  the  receiving 
vessel,  to  complete  the  emptying.  Apparatus  must  be  sufficiently 
clean  to  permit  uniform  wetting  of  the  surface.  (For  the  prepara- 
tion of  cleaning  solution,  see  the  Appendix.) 

Errors  Due  to  Parallax.  In  all  apparatus  in  which  the  volume 
is.  limited  by  a  meniscus,  the  reading  or  setting  is  made,  when 
possible,  on  the  lowest  point  of  the  meniscus.  Since  this  point 
lies  at  the  center  of  the  tube,  the  error  from  parallax  must  be 
avoided.  The  method  of  doing  this  is  to  support  the  tube  so 
that  its  main  axis  is  vertical,  and  to  hold  the  eye  at  such  a  level 
that  the  line  of  sight  makes  an  angle  of  90°  with  the  axis. 

In  order  that  the  lowest  point  of  the  meniscus  may  be  ob- 
served, it  is  well  to  place  a  screen  of  some  dark  material  im- 
mediately below  the  meniscus,  which  renders  the  profile  of  the 
meniscus  dark  and  clearly  visible  against  a  light  background. 
A  convenient  device  for  this  purpose  is  a  collar-shaped  section 
of  black  rubber  tubing,  cut  open  at  one  side  and  of  such  a  size 
as  to  clasp  the  tube  firmly. 

Errors  Due  to  Variations  in  Temperature.  The  volume  occu- 
pied by  a  given  weight  of  water,  as  well  as  the  capacity  of  the 
measuring  vessel,  is  dependent  upon  the  temperature ;  and  the 
error  involved  in  the  measurement  of  the  volume  of  a  given 
mass  of  water,  at  any  other  temperature  than  the  standard  one, 
is  due  to  the  joint  effect  of  the  changed  capacity  of  the  vessel 
and  the  changed  volume  of  the  liquid. 

The  coefficient  of  cubical  expension  of  ordinary  glass  may  be 
taken  to  be  0.000025 ;  but  the  volume  change  of  the  water  is 
much  greater  than  that  of  the  glass  measuring  vessel,  and  also 
much  less  uniform  from  degree  to  degree.  The  factors  by 


INTRODUCTION 


47 


which  a  volume  of  water,  measured  at  temperatures  ranging 
from  10-29°  m  a  vessel  calibrated  for  20°  must  be  multiplied 
in  order  to  obtain  the  true  volume  occupied  by  the  liquid  at 
20°,  are  given  in  the  following  table : 


TEMPERATURE  or  THE  WATER 

"""\^UNITS 
TENS^^-^ 

0 

* 

2 

3 

[4 

1 

i  .001  24 

I.OOII7 

1  .00109 

I.OOIOO 

1.00089 

2 

I.OOOOO 

0.99981 

0.99961 

0.99941 

0.99919 

TEMPERATURE  OF  THE  WATER 

"~^\UNITS 
TENS^^\_ 

5 

6 

7 

8 

9 

1 

1.00077 

1.00064 

I.OOO49 

1.00034 

I.OOOlS 

2 

0.99896 

0.99873 

0.99848 

0.99822 

0.99798 

If  the  prevailing  temperature  does  not  differ  by  more  than, 
say,  3°  from  the  standard,  this  correction  may  ordinarily  be 
omitted.  In  the  case  of  solutions  of  0.2  N  concentration,  or  less, 
the  corrections  differ  so  little  from  those  for  pure  water  that  the 
factors  given  in  the  table  may  be  used  without  appreciable  error. 

In  order  to  illustrate  the  use  of  such  factors,  let  us  suppose 
that,  in  the  standardization  of  a  solution,  a  burette  graduated 
correctly  for  20°  is  used  at  an  actual  temperature  of  27°,  and 
that  the  indicated  volume  of  solution  withdrawn  is  28.75  cc- 
Then  the  true  volume  at  20°  of  this  quantity  of  liquid  is  28.75  X 
0.99848  =  28.70  cc.  And,  if  a  determination  is  later  made  with 
this  solution  at,  say,  17°,  and  the  indicated  volume  used  is 
28.68  cc.,  then  the  true  volume  at  20°  is  28.68  X  1.00049  =  28.70  cc. 
That  is  to  say,  the  same  quantity  of  reagent  is  contained  in  an 
apparent  volume  of  28.75  cc-  at  27°>  or  in  an  apparent  volume 
of  28.68  cc.  at  17°,  as  is  contained  in  an  actual  volume  of  28.70  cc. 
of  the  solution  at  200.1 

1  At  27°  the  actual  volume  of  the  liquid  measured  is  a  shade  greater  than  28.75  cc-» 
and  at  17°  it  is  a  shade  less  than  28.68  cc.  The  measuring  vessel  is  larger  at  27°, 
and  smaller  at  17°,  than  it  is  at  20°. 


48     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Errors  due  to  Different  Units  of  Volume  Employed.  Unfor- 
tunately, a  number  of  different  "  liters  "  have  been  suggested 
for  use  in  volumetric  analysis.  The  normal  liter,  that  is,  the 
volume  occupied  by  a  kilogram  of  water,  weighed  in  a  vacuum 
and  measured  at  4°,  would  manifestly  be  out  of  the  question  if 
it  had  to  be  determined  in  that  way.  The  so-called  "  Mohr 
liter  "  is  the  volume  occupied  by  a  kilogram  of  water  when 
weighed  in  the  air  with  brass  weights  at  a  temperature  of  17.5° ; 
but  this  volume  varies  with  the  atmospheric  conditions.  Other 
"  liters  "  involving  measurements  at  15°,  15.5°,  or  20°  have  been 
suggested  by  various  chemists. 

It  matters  little,  in  analytical  work,  which  liter  is  adopted, 
but  it  is  of  the  greatest  importance  to  have  the  pipettes,  burettes, 
and  measuring  flasks  rigorously  consistent  with  one  another.  This 
matter  is  serious  enough  to  require  especial  emphasis,  since 
apparatus,  if  not  specifically  ordered,  may  be  supplied  by  dealers, 
at  different  times,  graduated  according  to  different  systems ;  and 
mixed  graduations  may  thus  come  into  the  hands  of  an  individual 
analyst.  As  an  example  of  the  magnitude  of  the  errors  which 
might  thus  be  introduced,  it  should  be  noted  that  the  normal 
liter  is  related  to  the  Mohr  liter  as  1000 :  1002.3. 

Much  of  the  graduated  apparatus  on  the  market  bears  no 
mark  by  means  of  which  the  unit  of  volume  represented  can  be 
recognized,  and  even  when  this  is  clearly  designated  the  per- 
centage error  represented  may  be  large.  It  is  not  advisable, 
therefore,  to  use  any  piece  of  graduated  apparatus,  unless  its 
actual  value  is  well  known. 

Owing  to  the  great  difficulty  in  measuring  directly  the  re- 
lation between  cubic  capacity  and  the  unit  of  length,  the  Inter- 
national Committee  of  Weights  and  Measures  defines  the  liter 
as  "  the  volume  occupied  by  the  mass  of  one  kilogram  of  pure 
water  at  its  maximum  density  under  normal  atmospheric  pres- 
sure."  This  is  almost  exactly  1000  cc.1  and  for  all  practical  pur- 
poses may  be  regarded  as  such. 

1  About  1000.029  cc. 


INTRODUCTION 


49 


It  is  now  customary  to  use  this  true  liter  as  the  standard, 
but  of  course  it  is  out  of  the  question  to  weigh  a  kilogram  of 
water  at  4°  in  a  vacuum ;  some  convenient  temperature  — 
preferably  the  average  working  temperature  of  the  laboratory 
—  must  be  selected,  and  the  necessary  corrections  made.  If  a 
liter  flask  is  marked  correctly  at  20°,  this  means  that  at  20°  it 
will  contain  a  mass  of  water  (998.234  g.)  which  occupies  a  volume 
equal  to  that  occupied  by  1000  grams  of  pure  water  at  4°.  This 
quantity  of  water,  if  weighed  with  brass  weights  in  air  of  mean 
humidity,  at  20°  and  760  mm.,  has  an  apparent  weight  of  997.18 
grams. 

The  Calibration  of  Volumetric  Apparatus.  The  weight  of 
brass  (brass  weights)  which  will  be  required  to  counterbalance 
one  liter  of  pure  water  must  be  calculated  from  the  temperature 
of  the  water  and  the  density  of  the  air.  The  following  table 
indicates,  for  temperatures  of  the  water  (and  room)  ranging  from 
15-29°,  how  many  milligrams  less  than  1000  grams  a  quantity  of 
water  will  weigh  which  is  sufficient  to  fill  to  the  mark  a  i -liter  flask 
correctly  calibrated  for  20°,  the  weighing  being  carried  out  in  air 
of  50  per  cent  humidity  at  760  mm.  pressure  (unreduced).1 


>NXNUNITS 
TENS\^ 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1 

— 

— 

'  — 

— 

— 

1950 

2IOO 

2260 

2440 

2630 

2 

2820 

3030 

3240 

3470 

3710 

3960 

42IO 

4480 

4760 

2620 

If  such  a  flask  is  filled  to  the  mark  with  water  of  22.4°,  for 
example,  the  water  will  under  the  conditions  of  the  table  require 
a  counterpoise  of  1000-3.332  =  996.668  grams. 

The  determination  of  the  capacity  of  a  measuring  flask  is 
carried  out  by  weighing  the  water  contained  in  it,  while  the 
volume  of  water  delivered  by  a  burette  or  pipette  is  determined 
by  weighing  this  water  after  its  delivery  into  another  vessel. 

1  These  values  depend  upon  the  specific  gravities  of  the  water  and  the  (brass) 
weights,  the  density  of  the  air,  and  the  coefficient  of  cubical  expansion  of  the  glass 
vessel. 


50     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  temperature  of  the  water  should  be  taken  before  and  after 
the  experiment ;  it  is  important  that  it  should  be  the  same  as 
the  room  temperature  at  the  time  of  the  weighing. 

In  the  calibration  of  a  flask,  the  dry  flask  is  placed  upon  the 
right-hand  pan  of  the  balance,  together  with  the  nominal  weight 
of  its  capacity,  i.e.  with  as  many  grams  as  it  is  supposed  to  con- 
tain cubic  centimeters,  and  then  tare  material  is  placed  upon 
the  left-hand  pan  until  the  balance  is  brought  into  equilibrium. 
The  weights  are  then  removed  from  the  right-hand  pan,  the 
flask  is  filled  to  the  mark  with  water,  and  weights  are  added 
until  the  balance  is  again  in  equilibrium.  The  nominal  capacity 
weight,  minus  the  additional  weight  which  is  required  upon  the 
right-hand  pan  in  order  to  reestablish  equilibrium,  is  equal  to 
the  weight  of  the  water  in  the  flask. 

In  the  case  of  burettes  and  pipettes,  a  covered  beaker  is  placed 
upon  the  right-hand  pan,  together  with  the  nominal  weight  in 
grams  of  the  volume  to  be  delivered,  after  which  the  balance 
is  brought  into  equilibrium  by  the  addition  of  tare  material 
to  the  left-hand  pan.  The  water  is  then  allowed  to  run  into  the 
beaker,  which  is  replaced  upon  the  right-hand  pan.  The  subse- 
quent procedure  is  the  same  as  that  described  above. 

The  difference  between  the  additional  weight  required  to 
reestablish  equilibrium  and  that  calculated  from  the  above 
table  indicates  directly  the  error  of  the  vessel.  If,  for  example, 
it  be  found  necessary,  in  order  to  reestablish  equilibrium  in  the 
case  of  a  5oo-cc.  flask  filled  to  the  mark  with  water  of  22.4°, 
to  add  1.832  g.,  instead  of  the  calculated  1.666  g.,  then  it  follows 
that  the  vessel  is  0.166  cc.  too  small.  On  the  other  hand,  in  test- 
ing the  25  cc.  segment  of  a  30-cc.  burette,  at  a  temperature  of  17°, 
the  additional  weight  required  on  the  right-hand  pan  should  be 
(25  X 2 260) -f- 1000  =  57  mg. ;  if,  instead  of  this,  it  be  found  that 
an  additional  weight  of  15  mg.  is  required  on  the  left-hand  pan, 
then  the  25  cc.  segment  is  57  mg.  — (  — 15  mg.)  =0.072  cc.  too 
large.1 

1  This  error  is  too  large  to  be  tolerated. 


INTRODUCTION  51 

In  the  calibration  and  use  of  burettes,  the  liquid  should  in 
general  be  allowed  to  flow  from  the  zero  mark  to  some  second 
level  in  the  burette. 

D.    THE  PREPARATION  OF  SAMPLES  FOR  ANALYSIS 

It  is  not  easy  to  give  general  rules  for  the  preparation  of  sub- 
stances for  analysis,  because  it  is  necessary  to  proceed  differ- 
ently in  different  cases.  In  all  cases,  however,  the  samples 
should  promptly  be  transferred  to  tightly  stoppered  bottles  or 
weighing  tubes. 

In  technical  analyses,  for  the  purpose  of  determining  the 
commercial  value  of  an  article,  or  of  controlling  processes  of 
manufacture,  materials  must  be  analyzed  as  they  are.  But, 
in  every  case,  especial  care  should  be  taken  to  make  up  a  sample 
which  will  represent  as  nearly  as  possible  the  average  composition 
of  the  whole  lot. 

If,  on  the  other  hand,  it  is  desired  to  determine  the  atomic 
composition  of  a  compound,  it  is  necessary  to  select  or  prepare 
pure  material  for  analysis.  This  may  seem  simpler  than  it 
really  is.  Many  compounds  absorb  or  give  up  moisture  upon 
exposure  to  the  air,  and  their  treatment  should  vary  with  their 
nature,  as  illustrated  in  the  following  cases.  Salts  such  has 
Na2SO4 . 10  H2O  and  Na2CO3 .  10  H2O,  which  effloresce  in 
ordinary  air,  may  be  dried,  after  recrystallization,  by  strongly 
pressing  the  powdered  material  between  several  layers  of  filter 
paper,  the  paper  being  renewed  as  long  as  moisture  continues 
to  be  taken  up ;  MgSO4 .  7  H2O  and  NaKC4H4O6 .  4  H20, 
which  do  not  lose  water  of  constitution  in  ordinary  air,  may  be 
spread  out  upon  filter  paper,  covered  with  another  sheet,  and 
allowed  to  dry  at  the  ordinary  temperature.  Compounds  such 
as  HFe(SO4)2 .  4  H2O  and  CaC4H4O6 .  H2O,  which  do  not 
effloresce  in  artificially  dried  air,  but  which  undergo  chemical 
change  at  100°,  may  be  conveniently  dried  in  a  desiccator,  over 
calcium  chloride.  Substances,  as  KHC4H4O6,  sugar,  etc.,  which 
readily  give  up  hygroscopic  moisture  at  100°,  without  other 


52     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

alteration,  are  best  dried  in  an  oven  at  that  temperature ;  while 
K2PtCl6,  which  retains  moisture,  or  dries  only  slowly  at  100°, 
but  which  decomposes  below  a  red  heat,  should  be  dried  in  an 
oven  at,  say,  130°.  Finally,  substances  such  as  NaCl,  Na2S04, 
etc.  may  be  given  a  preliminary  drying,  in  a  covered  vessel,  at 
130°,  or  higher,  to  prevent  decrepitation,  and  then  be  ignited, 
more  or  less  strongly,  depending  upon  their  nature.  In  every 
case,  the  sample  should  be  dried,  without  decomposition,  to 
constant  weight. 

Substances  used  in  testing  the  accuracy  of  analytical  pro- 
cesses, or  in  standardizing  volumetric  solutions,  must  also  be 
extremely  pure.  In  fact,  compounds  are  generally  favored 
which  are  non-hygroscopic,  and  which  may  readily  be  prepared 
in  a  pure  condition;  if  possible,  it  is  well  to  select  compounds 
which  normally  do  not  contain  water  of  crystallization.  Many 
salts  can  be  obtained  sufficiently  pure  in  the  market ;  but  their 
purity  should  never  be  accepted  on  faith.  If  tests  indicate  the 
presence  of  impurity,  and,  often,  if  the  salt  contains  water  of 
crystallization,  the  material  should  be  recrystallized. 

For  this  purpose,  a  convenient  weight  of  the  salt  is  dissolved 
in  the  least  possible  quantity  of  hot  water,  using  a  quantity  of 
water  not  quite  sufficient  to  dissolve  the  whole  lot;  the  hot 
solution  is  poured  into  a  fluted  filter,  held  in  a  stemless  funnel, 
and  the  filtrate  is  received  with  continuous  stirring  in  a  beaker, 
which  itself  is  immersed  in  cold  water,  in  a  larger  vessel.  The 
rapid  cooling  and  constant  stirring  cause  the  formation  of  a 
fine  crystalline  powder,  which  is  almost  free  from  inclosed 
mother-liquor.  The  crystalline  powder  is  filtered  off  in  a  funnel 
containing  a  perforated  platinum  cone,  the  adhering  mother- 
liquor  being  removed  by  suction  or  in  a  centrifuge.  Two  such 
recrystallizations  will  nearly  always  suffice.  According  to  the 
nature  of  the  substance,  it  is  dried  in  the  air  at  a  specific  tem- 
perature, or  in  a  desiccator,  to  constant  weight. 

Concerning  the  preparation  of  samples  for  analysis  by  beginners 
in  quantitative  analysis,  the  reader  should  consult  the  appendix. 


PART    II 

GRAVIMETRIC  ANALYSIS 
EXERCISES  WITH  THE  BALANCE 

Before  beginning  work  at  the  balance,  read  carefully  the 
rules  given  on  pp.  9-11  of  Part  I,  and  observe  them  always. 

Determination  of  the  Zero-Point.  Determine  the  zero-point 
of  the  unloaded  balance,  according  to  the  method  on  p.  n.  If 
the  zero-point  found  is  not  more  than  one  division  from  the 
center  of  the  scale,  the  balance  may  be  used  by  the  student; 
otherwise  it  will  be  adjusted  by  an  instructor,  upon  request. 
The  beginner  should  not  attempt  this  adjustment. 

Determination  of  the  Weight  of  an  Object.  Clean  two  porce- 
lain crucibles,  rinse  them  with  distilled  water,  and  allow  them 
to  drain.  Place  each  crucible  upon  a  pipes  tern  triangle,  sup- 
ported upon  a  tripod,  and  heat  with  the  colorless  flame  of  a 
Bunsen  or  Tyrill  burner,  —  gently  at  first,  and  then  to  a  red 
heat.  Allow  the  crucibles  to  cool  off  somewhat,  but  while  still 
warm,  place  them  in  a  desiccator,  using  the  crucible  tongs. 
(After  any  piece  of  apparatus  has  been  cleaned  and  ignited  for 
weighing,  it  must  never  be  handled  with  the  fingers  before  the  weight 
is  taken.}  Allow  the  crucibles  to  cool  in  the  desiccator  for  at 
least  20  minutes. 

Now  with  the  crucible  tongs  place  a  crucible  on  the  left-hand 
pan  of  the  balance,  and,  by  means  of  the  forceps  in  the  weight 
box,  place  weights  on  the  right-hand  pan  until  they  balance  the 
crucible  to  within  0.005  g-  Begin  with  a  weight  which  you 
think  will  approximately  balance  the  object,  lower  the  balance 
beam,  and  gently  release  the  pan  supports.  It  will  then  be  seen 

53 


54     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

which  side  is  the  heavier.  Finally  adjust  the  rider,  so  that, 
when  the  beam  is  swinging  freely,  the  pointer  traverses  the  same 
number  of  divisions  on  either  side  of  the  observed  zero-point. 
Always  try  the  weights  in  the  order  in  which  they  occur  in  the 
box,  beginning  with  the  heavier  weights,  and  using  the  rider  for 
weights  smaller  than  5  or  10  milligrams,  according  to  the  num- 
ber of  large  divisions  on  the  beam. 

As  soon  as  the  object  appears  to  be  balanced,  raise  and  lower 
the  beam,  and  make  another  observation.  Read  the  weight  of 
the  crucible  by  noting  in  order  the  vacant  spaces  in  the  box,  begin- 
ning with  the  largest  missing  weight;  and  check  this  reading  as  the 
weights  are  returned  to  the  box .  Be  sure  also  to  note  the  weight 
recorded  by  the  rider,  and  then  lift  it  from  the  beam.  Always 
record  the  weight,  in  pencil  as  it  is  first  read,  and  in  ink  after  it 
has  been  checked,  and  always  in  the  record  book,  —  never  on  a 
loose  sheet  of  paper. 

In  this  manner,  weigh  the  two  crucibles  separately,  and  then 
weigh  them  together,  entering  all  three  results  in  the  notebook. 
(In  connection  with  the  keeping  of  records,  see  the  remarks  on 
p.  4.)  The  sum  of  the  separate  weights  should  agree  closely 
with  the  result  obtained  upon  weighing  both  crucibles  together, 
—  within,  say,  0.0002  g. 

THE  DETERMINATION  OF  CHLORINE  IN  A  SOLUBLE 
CHLORIDE 

The  sample  may  be  pure  sodium  chloride,  or  it  may  be  an 
artificially  prepared  mixture  of  sodium  chloride  and  sodium 
carbonate. 

Method.  The  aqueous  solution  of  the  chloride  is  acidified 
with  nitric  acid  and  treated  with  silver  nitrate  in  excess.  The 
chlorine  is  quantitatively  precipitated  as  silver  chloride,  which 
is  filtered  off,  washed,  dried,  and  weighed.  Other  acids  which 
yield  silver  salts  insoluble  in  nitric  acid  must  of  course  be  absent. 

A.  A  Paper  Filter  is  Used:  Procedure.  Carefully  clean  the 
weighing  tube  containing  the  sample,  without  handling  it  di- 


GRAVIMETRIC  ANALYSIS  55 

rectly  with  the  fingers,  and  weigh  it  accurately  to  a  tenth  of  a 
milligram.  Record  the  weight  at  once  in  the  notebook.  Hold 
the  tube  over  a  clean  300-0:.  beaker  (plainly  labelled  "  I  "), 
remove  the  stopper,  allowing  no  particles  to  fall  from  it  or  from 
the  tube  elsewhere  than  into  the  beaker,  and  carefully  pour 
into  the  beaker  from  0.2  to  0.3  g.  of  the  sample.  Replace  the 
stopper  in  the  tube,  weigh  accurately,  and  record  the  weight  in 
the  notebook.  The  difference  of  these  two  weights  is  the  weight 
of  the  portion  taken  for  analysis.  Weigh  a  second  portion  of 
0.2-0.3  g.  into  another  beaker  (labelled  "  II  "),  entering  the 
weights  and  their  difference  in  the  notebook,  as  before. 

Dissolve  each  portion  in  about  150  cc.  of  distilled  water,  and 
acidify  the  solutions  with  nitric  acid,  adding  the  acid  slowly 
and  with  stirring,  until  a  strip  of  blue  litmus  paper  shows  an 
acid  reaction  when  moistened  by  means  of  the  wet  stirring  rod 
with  the  least  possible  quantity  of  the  liquid.  Assuming  the 
sample  to  be  pure  sodium  chloride,  calculate  the  volume  of 
silver  nitrate  solution  required  in  each  case  to  effect  complete 
precipitation  (for  the  strength  of  the  reagents,  see  the  Appendix), 
and  add  slowly  and  with  stirring  about  5  cc.  more  than  that 
amount.  Cover  the  beaker  with  a  watch  glass,  and  heat  the 
solution  gradually  to  boiling,  with  occasional  stirring.  Con- 
tinue the  heating  and  stirring  until  the  precipitate  coagulates 
and  the  supernatant  liquid  is  clear.  The  beaker  should  be  kept 
away  from  direct  sunlight,  and  the  heating  and  stirring  should 
be  so  conducted  as  to  avoid  any  possibility  of  loss.  Finally, 
add  to  the  clear  liquid  a  drop  or  two  of  silver  nitrate  solution, 
to  test  for  complete  precipitation;  if  a  precipitate  forms,  add 
5  cc.  more,  and  test  again. 

Prepare  two  ashless  filters  (9  cm.  in  diameter),  according  to 
the  directions  given  on  p.  30,  and  decant  the  hot  liquid  through 
the  filter  in  each  case,  leaving  the  precipitate  as  far  as  possible 
in  the  beaker.  Unless  the  filtrates  are  perfectly  clear  and  free 
from  particles  of  silver  chloride,  they  must  be  refiltered  through 
the  same  filters.  If  they  are  perfectly  clear,  and  the  tests  show 


56     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

complete  precipitation,  pour  them  into  one  of  the  laboratory 
receptacles  for  "  Silver  Residues,"  wash  the  beakers  with  tap 
water  and  then  with  distilled  water,  and  replace  them  under 
the  funnels.  Now  wash  the  precipitates  twice  by  decantation 
with  10  cc.  portions  of  hot  water,  acidified  with  a  drop  or  two 
of  nitric  acid,  pouring  the  washings  through  the  filters,  and  finally 
transfer  each  precipitate  to  the  corresponding  filter  by  means  of 
a  stream  of  hot  water  from  the  wash  bottle,  loosening  the  adher- 
ing particles  with  the  aid  of  a  " policeman"  (see  p.  32).  Wash 
the  filters  and  precipitates  with  hot  water  until  3  cc.  of  the  wash- 
ings show  no  cloudiness  or  opalescence  with  one  drop  of  dilute  hy- 
drochloric acid.  After  allowing  the  filters  to  drain,  cover  each 
funnel  with  an  ordinary  filter  paper,  crimping  the  edges  of  the 
paper  over  the  sides  of  the  funnel.  The  funnels,  properly 
numbered,  and  labeled  with  the  student's  name  and  desk  num- 
ber, should  then  be  placed  in  a  drying  oven,  at  a  temperature  of 
9o°-ioo°,  and  left  there  until  completely  dry. 

Now,  in  the  case  of  each  precipitate,  open  the  filter  over  a 
piece  of  smooth  glazed  paper,  about  six  inches  in  diameter,  and, 
by  finally  rubbing  the  sides  of  the  filter  gently  together,  transfer 
the  precipitate  as  completely  as  possible  to  the  center  of  the  glazed 
paper.  Be  careful  not  to  rub  off  any  appreciable  quantity  of 
the  paper,  nor  to  lose  any  of  the  silver  chloride  in  the  form  of 
dust.  Cover  the  precipitate  on  the  paper  with  an  inverted  fun- 
nel or  watch  glass,  to  protect  it  from  dust  and  air  currents. 

Carefully  refold  the  paper,  flat,  bend  the  top  over,  and  roll  the 
paper  into  a  small  bundle ;  then  place  it  in  a  weighed  porcelain 
crucible.  Place  the  crucible  upon  a  triangle,  incline  it  about  45°, 
and  ignite  gently  until  the  volatile  products  are  expelled  from  the 
paper.  Then,  with  the  flame  well  at  the  base  of  the  inclined 
crucible,  ignite  strongly  until  all  the  carbon  is  consumed.  (See 
Part  I,  p.  38.)  Allow  the  crucible  to  cool,  add  two  drops  of 
6-normal  nitric  acid  and  one  of  hydrochloric  acid,  and  heat  with 
the  greatest  caution,  to  avoid  spattering,  until  the  acids  are  ex- 
pelled. Transfer  the  bulk  of  the  precipitate  quantitatively  from 


GRAVIMETRIC  ANALYSIS  57 

the  glazed  paper  to  the  cooled  crucible,  placing  the  latter  on  a 
second  piece  of  glazed  paper  and  brushing  the  precipitate  into 
it,  with  a  small  camel's  hair  brush. 

Moisten  the  precipitate  with  two  drops  of  nitric,  and  one 
drop  of  hydrochloric  acid,  carefully  expel  the  acids,  and  then 
gradually  raise  the  temperature  until  the  salt  just  begins  to  fuse. 
Allow  the  crucible  to  cool  in  a  desiccator,  and  weigh  it.  Repeat 
the  heating,  without  the  addition  of  acids,  and  weigh  the  cooled 
crucible.  The  heating  and  weighing  must  be  repeated  until 
the  weight  is  constant  within  0.2  mg.  after  two  consecutive 
heatings.  From  the  weight  of  silver  chloride  obtained  in  each 
case,  calculate  the  percentage  of  chlorine  in  the  sample. 

Finally,  place  the  ignited  precipitates  in  one  of  the  laboratory 
receptacles  for  silver  residues.  The  chloride  which  adheres  to 
the  crucible  may  be  loosened  by  covering  it  with  dilute  sulphuric 
acid  and  adding  a  small  quantity  of  granulated  zinc. 

NOTES.  —  i.  The  solution  is  acidified  with  nitric  acid,  before  precipita- 
tion with  silver  nitrate,  to  prevent  the  precipitation  of  substances  such  as 
silver  oxide,  carbonate,  phosphate,  etc.,  which  are  insoluble  in  water  but 
soluble  in  nitric  acid.  The  acid  also  helps  to  coagulate  the  precipitate.  A 
large  excess  of  the  acid  is  to  be  avoided,  since  it  would  slightly  increase  the 
solubility  of  the  precipitate. 

2.  It  is  safer  not  to  boil  the  acidified  solution  until  after  silver  nitrate 
has  been  added  in  excess,  since  otherwise  a  slight  amount  of  chlorine  might 
be  set  free  by  the  nitric  acid  and  lost  by  volatilization.    The  presence  of  an 
excess  of  silver  nitrate  can  easily  be  recognized  at  the  time  of  its  addition, 
by  the  increased  readiness  with  which  the  precipitate  coagulates  and  settles. 

3.  The  precipitate  should  not  be  exposed  to  strong  sunlight,  since  by  its 
action  a  slight  amount  of  chlorine  is  set  free.     The  superficial  alteration 
which  the  chloride  undergoes  in  diffused  daylight  may  readily  be  counter- 
acted by  the  treatment  with  nitric  and  hydrochloric  acids ;  but  the  loss  in 
weight  due  to  this  cause  is  really  too  insignificant  to  affect  the  accuracy  of 
the  determination. 

4.  The  precipitate  and  filter  are  washed  with  water  to  remove  the  non- 
volatile nitrates  of  silver  and  sodium,  as  well  as  any  other  soluble  impurities. 
It  may  be  assumed  that  these  are  all  removed  as  soon  as  3  cc.  of  the  wash- 
ings give  no  cloudiness  with  a  drop  of  hydrochloric  acid.     Only  a  single 
drop  should  be  added,  because  silver  chloride  is  somewhat  soluble  in  this 


58     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

acid.  The  wash  water  should  be  hot  in  order  to  prevent  the  precipitate 
from  going  into  colloidal  solution ;  it  is  still  safer  to  acidify  the  water  slightly 
with  nitric  acid. 

5.  The  bulk  of  the  precipitate  must  be  separated  from  the  filter,  be- 
cause the  burning  organic  matter  would  reduce  a  considerable  quantity  of 
the  precipitate  to  metallic  silver,  and  its  complete  reconversion  to  the 
chloride  within  the  crucible,  by  means  of  acids,  would  be  uncertain.     The 
small  quantity  which  adheres  to  the  filter,  and  which  is  more  or  less  com- 
pletely reduced  during  ignition,  is  easily  reconverted  to  chloride  by  the 
treatment  with  nitric  and  hydrochloric  acids. 

6.  Silver  chloride  should  not  be  heated  to  complete  fusion,  since  a  slight 
loss  by  volatilization  might  take  place.    The  temperature  of  fusion  is 
sufficient  to  completely  remove  adsorbed  water  and  acids,  but  it  is  not 
always  sufficient  to  destroy  filter  shreds.    Although  these  would  probably 
be  completely  oxidized  by  the  nitro-hydrochloric  acid  and  subsequent 
ignition,  they  should  not  be  allowed  to  contaminate  the  precipitate. 

7.  The  ignited  precipitate  of  silver  chloride,  as  well  as  the  filtrates  which 
contain  an  excess  of  silver  nitrate,  should  be  placed  in  the  receptacles  for 
silver  residues ;  the  silver  can  easily  be  recovered. 

Assuming  that,  on  the  average,  duplicate  determinations  require  55  cc. 
of  o.2-normal  silver  nitrate  solution,  the  residues  returned  in  the  case  of 
each  student  should  contain  about  1.2  g.  of  metallic  silver.  Taking  into 
account  the  number  of  analyses  which  have  to  be  repeated,  a  class  of  one 
hundred  students  will  usually  return  in  residues  at  least  150  g.  (about  5  oz.) 
of  metallic  silver. 

8.  Silver  chloride  is  almost  insoluble  in  water.    The  solubility  varies 
somewhat  with  its  physical  condition,  and  is  about  1.12  mg.  per  liter  at 
20°  C.     Owing  to  the  common-ion  effect,  the  solubility  is  still  less  in  a  very 
dilute  solution  of  silver  nitrate  or  of  hydrochloric  acid.     In  hot  water  the 
salt  is  more  soluble,  21.8  mg.  per  liter  at  100°  C. ;  but,  fortunately,  the 
speed  of  solution  is  so  slow  that  the  precipitate  may  be  thoroughly  washed 
with  hot  water,  without  undue  error  from  this  cause. 

Silver  ion  has  a  great  tendency  to  enter  into  the  formation  of  complex  ions, 
as  [Ag(NH3)2]+,  [Ag(CN)2]-,  [AgS2O3]-,  etc.,  and  the  chloride  is  therefore 
readily  soluble  in  aqueous  ammonia,  in  alkali  cyanide  solutions,  and  in 
sodium  thiosulphate  ("hypo")  solution.  Silver  chloride  is  also  soluble  in 
strong  hydrochloric  acid  and  in  other  chloride  solutions,  probably  with  the 
formation  of  a  complex  anion,  such  as  [AgQ2]~;  it  is  also  soluble  in  con- 
centrated silver  nitrate  solution,  and  in  strong  nitric  acid.  When  boiled 
with  concentrated  sulphuric  acid,  it  is  converted  into  silver  sulphate,  and 
by  zinc  and  dilute  sulphuric  acid  it  is  reduced  to  metallic  silver. 


GRAVIMETRIC  ANALYSIS  59 

B.  A  Gooch  Crucible  is  Used:  Procedure.  Weigh  out  two 
samples  of  the  substance,  of  about  0.25  g.  each,  and  convert  the 
chloride  into  silver  chloride  as  described  in  procedure  A .  Mean- 
while, prepare  two  Gooch  crucibles,  following  the  directions  on  pp. 
33-35  ;  and  finish  the  analysis  according  to  the  details  there  given. 

NOTE.  —  Bromides,  iodides,  cyanides,  sulphocyanates,  etc.,  as  well  as 
silver  itself,  may  be  determined  in  a  similar  manner,  with  the  use  of  Gooch 
crucibles.  Chlorates,  etc.,  may  be  determined  by  first  reducing  them  to 
chlorides,  and  then  precipitating  with  silver  nitrate. 

For  the  determination  of  these  substances  when  two  or  more  of  them  are 
present  in  the  same  sample,  the  student  is  referred  to  the  larger  works  on 
quantitative  analysis.  (But  see  also  Part  IV,  Problems  37,  38,  90,  91, 
and  93.) 

THE    DETERMINATION   OF   IRON   AND    OF   SULPHUR   IN   A 
SOLUBLE   SULPHATE  OF  IRON 

The  sample  may  be  pure  ferrous  ammonium  sulphate,  pure 
ferric  alum,  or  an  artificially  prepared  mixture  of  anhydrous 
ferric  sulphate,  sodium  carbonate,  and  potassium  sulphate. 
This  mixture  is  readily  soluble  in  dilute  hydrochloric  acid. 

Method.  The  sample  is  dissolved  in  water,  with  the  addition 
of  hydrochloric  acid,  after  which  the  iron  is  oxidized  to  the  ferric 
condition,  unless  it  is  already  wholly  present  in  that  state.  The 
iron  is  then  separated,  by  double  precipitation  with  ammonium 
hydroxide,  as  ferric  hydroxide.  The  precipitate  is  ignited,  and 
weighed  as  ferric  oxide. 

From  the  combined  filtrates  and  washings,  which  must  have 
a  large  volume,  and  which  must  be  free  from  nitrates,  etc.,  the 
sulphate  is  precipitated  by  means  of  a  dilute  solution  of  barium 
chloride.  The  precipitate  is  ignited,  and  weighed  as  barium 
sulphate. 

A.  Procedure  for  the  Determination  of  Iron.  Weigh  out 
into  dry  200  cc.  beakers  two  portions  of  about  one  gram  each, 
and  add  to  each  portion  50  cc.  of  water  and  10  cc.  of  6-normal 
hydrochloric  acid,  keeping  the  beakers  covered  with  watch- 


60     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

glasses  to  prevent  loss  by  effervescence.  Treat  each  solution  as 
follows :  Heat  to  boiling,  and  add,  drop  by  drop,  nitric  acid  (sp. 
gr.,  I.42),1  until  the  brown  coloration  at  first  imparted  to  the 
liquid  gives  place  to  a  yellow  or  red.  (Note  the  volume  of  nitric 
acid  which  is  added ;  i  cc.  will  always  be  found  sufficient.)  Boil 
for  three  minutes,  and  pour  the  solution,  with  stirring,  into  an 
excess  of  ammonia  which  has  been  diluted  with  water  to  a  volume 
of  200  cc.  (For  this  purpose,  calculate  the  volume  of  6-normal 
ammonium  hydroxide  required  to  neutralize  the  acids  added, 
and  use  5  cc.  in  addition.)  Heat  the  mixture  to  boiling,  and 
allow  the  precipitate  to  settle.  Decant  the  boiling-hot,  clear 
liquid  through  an  ashless  filter  (9  cm.  in  diameter),  leaving  the 
precipitate  as  far  as  possible  in  the  beaker,  and  wash  twice  by 
decantation  with  50  cc.  portions  of  very  hot  water,  still  leaving 
the  precipitate  in  the  beaker.  (At  once  neutralize  the  filtrate 
and  washings  with  hydrochloric  acid,  and  reserve  for  the  sul- 
phate determination.  Their  evaporation,  as  directed  under  the  deter- 
mination of  sulphur,  should  be  begun  at  this  point.  See  p.  4.) 

Dissolve  the  precipitate  by  pouring  through  the  filter  a  boiling 
mixture  of  5  cc.  of  water  and  10  cc.  of  6-normal  hydrochloric  acid, 
adding  the  acid  in  small  portions,  and  collecting  the  filtrate  (and 
washings)  in  the  beaker  containing  the  bulk  of  the  precipitate, 
which  also  should  completely  dissolve.  After  thoroughly  washing 
the  filter,  tear  it  into  small  bits  and  add  these  to  the  ferric  chlo- 
ride solution.  Pour  this  mixture  into  an  excess  of  ammonia,  as 
before,  and  heat  to  boiling.  Filter  boiling-hot  through  a  fresh 
filter,  wash  the  precipitate  twice  by  decantation  with  hot  water, 
and  finally  transfer  it  to  the  filter ;  wash  continuously  with  hot 
water  until  3  cc.  of  the  washings  show  no  turbidity  when  treated 
with  a  drop  of  nitric  acid  and  one  of  silver  nitrate  solution. 
(The  combined  filtrate  and  washings  should  at  once  be  neu- 
tralized with  hydrochloric  acid  and  added  to  those  from  the  first 
precipitation ;  the  evaporation  should  be  allowed  to  continue.) 

Ignite  the  precipitate,  together  with  the  filter,  in  an  inclined 
1  Before  adding  nitric  acid,  see  Note  2. 


GRAVIMETRIC  ANALYSIS  '  6l 

platinum  or  porcelain  crucible.  After  the  volatile  matter  of 
the  filter  has  been  expelled,  raise  the  temperature  to  the  full 
heat  of  the  burner,  and,  with  the  flame  well  at  the  base  of  the 
inclined  crucible,  continue  the  heating  for  about  15  minutes. 
Cool  in  the  desiccator,  and  weigh.  Repeat  the  heating  until 
the  weight  is  constant  within  0.2  mg.  Report  the  percentage 
of  iron  in  the  sample. 

NOTES.  —  i.  In  ferrous  salt  solutions  the  iron  is  slowly  oxidized  by 
oxygen  from  the  air  (4  Fe+++02+2  H20  =  4  Fe++++4  OH~),  and,  unless 
the  solution  contains  free  acid  to  remove  the  OH~  ions,  the  iron  will  partially 
precipitate  in  the  form  of  a  basic  ferric  sulphate.  Moreover,  owing  to 
hydrolysis,  upon  boiling  an  aqueous  solution  of  ferric  sulphate  there  results 
a  partial  precipitation  of  the  iron,  again  in  the  form  of  a  basic  ferric  sulphate 
(e.g.  Fe2(SO4)3+2H20=2  Fe(OH)S04+H2SO4).  This  action,  however,  is 
prevented  by  the  presence  of  sufficient  hydrochloric  acid. 

2.  The  complete  oxidation  of  the  iron  is  necessary,  since  ferrous  iron  is 
not  quantitatively  precipitated  by  ammonia.  In  the  absence  of  air,  in- 
deed, ammonium  salts  are  capable  of  preventing  entirely  the  precipitation 
of  iron  from  ferrous  salt  solutions.  Ferric  iron,  in  the  absence  of  organic 
matter,  is  completely  precipitated  by  ammonia,  even  in  the  presence  of 
ammonium  salts. 

The  nitric  acid  oxidizes  the  iron  according  to  the  equation :  6  FeS04+ 
2  HN03+6  HC1=  2  Fe2(S04)3+2  FeCl3+2  NO+4  H20,  and  the  dark  color 
imparted  to  the  solution  is  due  to  the  union  of  the  nitric  oxide  with  ferrous 
salt  which  has  not  yet  been  oxidized,  to  form  an  unstable  nitroso-compound 
similar  to  that  formed  in  the  "ring- test"  for  nitrates.  The  nitric  oxide  is 
expelled  by  heat,  and  the  solution  finally  acquires  the  yellow  color  which  is 
characteristic  of  ferric  chloride  in  the  presence  of  hydrochloric  acid. 

To  insure  the  presence  of  iron  wholly  in  the  ferric  condition,  a  very  small 
quantity  of  the  oxidized  solution  should  be  tested  on  a  porcelain  plate  with 
a  drop  of  very  dilute,  freshly  prepared  potassium  ferricyanide  solution 
(a  piece  of  the  salt  the  size  of  a  pinhead,  in  20  cc.  of  water) .  If  the  solution 
has  a  volume  of  50  cc.,  and  an  ordinary  drop  a  volume  of  0.05  cc.,  then  the 
loss  of  the  latter  would  occasion  an  experimental  error  of  0.1%  in  the  iron 
(and  in  the  sulphur)  determination.  In  this  case,  therefore,  we  might  add 
a  couple  of  drops  of  the  solution  to  i  cc.  of  water  in  a  clean  watch  glass,  and 
use  one  drop  of  this  mixture  for  the  test ;  the  remainder  can  then  be  washed 
back  into  the  beaker  containing  the  bulk  of  the  solution.  The  error  is  thus 
reduced  to  0.01%,  which  is  negligible. 


62     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Much  time  can  often  be  saved  by  testing  in  this  way  the  solution  made 
from  a  very  small  portion  of  the  (un weighed)  original  sample,  for  ferrous 
iron ;  in  its  absence  the  addition  of  the  nitric  acid,  and  the  subsequent  evapo- 
ration to  dryness,  should  be  omitted. 

3.  If  ammonia  is  added  to  ferric  sulphate  solution,  the  ferric  hydroxide 
precipitate  is  apt  to  be  contaminated  with  sulphuric  acid,  in  the  form  of  a 
basic  ferric  sulphate.     A  gradual  neutralization  with  ammonia  is  almost 
sure  to  lead  to  the  separation  of  an  insoluble  basic  sulphate,  owing  to  a 
deficiency  of  hydroxide  ions.     If,  however,  the  iron  solution  is  added  with 
stirring  to  an  excess  of  aqueous  ammonia,  the  precipitate  obtained  will  be 
comparatively  free  from  basic  sulphate.     But,  since  even  here  the  basic 
salt  is  likely  to  be  present  in  small  quantity,  the  precipitate  is  redissolved 
and  the  solution  again  added  to  an  excess  of  dilute  ammonia. 

4.  To  avoid  errors  due  to  the  solvent  action  of  ammonium  hydroxide 
upon  glass,  the  precipitate  should  be  filtered  off  without  unnecessary  delay. 
It  is  for  this  reason,  also,  that  the  filtrates  and  washings  should  be  promptly 
neutralized    with    hydrochloric    acid.    The    ferric    hydroxide    precipitate 
should  under  no  circumstances  be  allowed  to  dry  before  the  washing  has 
been  completed;  it  would  be  sure  to  crack,  and  in  a  subsequent  washing 
the  wash  water  would  simply  run  through  the  crevices. 

5.  During  the  combustion  of  the  filter,  a  portion  of  the  precipitate  may 
be  reduced  to  FesO^  and  it  is  essential  that  any  of  this  substance  should 
be  oxidized  back  to  ferric  oxide.     Therefore,  during  the  ignition,  it  is  im- 
portant that  there  should  be  a  ready  access  of  air  to  the  precipitate.     For 
this  reason  it  is  directed  to  macerate  the  filter  with  the  solution  of  ferric 
chloride  before  the  second  precipitation ;  this  insures  a  very  porous  mass 
which  is  readily  reoxidized. 

6.  The  foregoing  method  may  be  used  for  the  gravimetric  determination 
of  aluminum  or  chromium,  with  the  additional  precaution  that  the  solution, 
before  it  is  filtered,  must  be  heated  until  but  a  very  slight  excess  of  ammonia 
remains,  the  hydroxides  of  these  metals  being  more  soluble  in  aqueous 
ammonia  than  ferric  hydroxide. 

If  it  is  desired  by  this  method  to  determine  the  chromium  in  an  alkali 
chromate,  the  latter  is  boiled  with  hydrochloric  acid  and  alcohol,  in  order 
to  reduce  the  chromium  to  the  trivalent  condition. 

(K2Cr207+3  C2H60+8  HC1=2  KC1+2  CrCl3+3  C2H40+7  H20). 

The  hydroxides  of  all  three  metals  are  precipitated  by  sodium  or  potas- 
sium hydroxide,  but  the  precipitates  are  always  contaminated  with  alkali. 
Furthermore,  aluminum  and  chromium  hydroxides  dissolve  readily  in  an 
excess  of  caustic  alkali  and  form  anions,  to  which  the  formulas  A102~  and 
CrO2-  are  usually  ascribed.  When  freshly  precipitated,  all  three  hydroxides 


GRAVIMETRIC  ANALYSIS  63 

dissolve  in  hydrochloric  acid ;  but  aluminum  hydroxide,  after  standing  for 
some  time,  dissolves  with  considerable  difficulty.  While  their  precipitation 
is  favored  by  the  presence  of  ammonium  salts  (coagulation),  it  is  entirely 
prevented  by  the  presence  of  sufficient  tartaric  acid  (formation  of  soluble  com- 
plexes) ;  citric  acid,  glycerol,  sugars,  etc.  resemble  tartaric  acid  in  this  respect. 
Upon  ignition,  each  hydroxide  yields  an  oxide  suitable  for  weighing  — 
Fe2O3,  A12O3,  Cr203.  Chromic  oxide,  however,  upon  ignition,  is  partially 
oxidized  to  Cr2(CrO4)3;  it  should  be  ignited  in  a  current  of  hydrogen  (G. 
Rothaug,  Zeitschrift  jur  anorganische  Chemie,  Vol.  84,  pp.  165-189  (1913)). 

B.  Procedure  for  the  Determination  of  Sulphur.  Evaporate 
to  dryness,  on  the  steam  bath,  the  combined  nitrates  and  wash- 
ings from  the  iron  determination ;  add  to  the  residue  10  cc.  of 
6-normal  hydrochloric  acid;  and  again  evaporate  to  dryness.1 
Dissolve  the  residue  in  100  cc.  of  water,  and  filter  the  solution  if 
it  is  not  perfectly  clear.  Transfer  the  solution  to  a  700  cc.  beaker, 
dilute  to  400  cc.,  and  then  add  1.5  cc.  of  6-normal  hydrochloric 
acid.  Heat  the  solution  to  boiling,  and,  with  stirring,  quickly 
pour  in  a  boiling  mixture  of  15  cc.  of  i-normal  barium  chloride 
solution  and  100  cc.  of  water.  Continue  the  boiling  and  stirring 
for  two  or  three  minutes ;  allow  the  precipitate  to  settle,  and,  at 
the  end  of  half  an  hour,  after  testing  for  complete  precipitation, 
decant  the  liquid  through  a  filter.  Substitute  a  clean  beaker  for 
the  beaker  containing  the  clear  filtrate,  wash  the  precipitate  by 
decantation  with  hot  water,  and  subsequently  upon  the  filter 
with  hot  water  until  3  cc.  of  the  washings  give  no  cloudiness  or 
opalescence  with  a  drop  of  silver  nitrate  solution  and  one  of 
nitric  acid.  The  precipitate  is  then  dried  and  ignited  to  constant 
weight,  with  the  flame  well  at  the  base  of  the  inclined  crucible. 
Report  the  percentage  of  S04  in  the  sample. 

NOTES.  —  i.  Barium  sulphate,  to  a  greater  degree  than  most  precipi- 
tates, tends  to  carry  down  other  salts  which  are  present  in  the  solution  from 
which  it  separates,  and  these  substances  cannot  be  removed  by  simply 
washing  the  precipitate.  This  is  especially  true  of  nitrates  and  chlorates, 

1  In  case  the  original  sample  contained  no  ferrous  iron,  and  the  addition  of  nitric 
acid  was  omitted,  it  is  sufficient  only  to  evaporate  the  neutralized  nitrates  and 
washings  to  about  300  cc.  This  liquid  is  then  transferred  to  the  beaker,  diluted 
to  about  400  cc.,  and  treated  further  as  described  in  the  procedure. 


64     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

and  of  the  salts  of  trivalent  metals,  such  as  iron,  chromium,  etc.  There- 
fore, if  nitric  acid  has  been  used  to  oxidize  the  iron,  it  must  be  completely 
removed  by  evaporation  with  a  large  excess  of  hydrochloric  acid. 

Iron  is  always  present  in  the  sulphate  precipitated  from  hot  solutions  in 
the  presence  of  ferric  salts,  and  the  precipitate  then  loses  sulphuric  acid 
upon  ignition,  and  gives  low  results  in  spite  of  its  iron  content.  Pure 
barium  sulphate  itself  is  not  decomposed  at  a  red  heat,  but  suffers  loss, 
probably  of  sulphur  trioxide,  at  a  temperature  above  900°. 

Barium  sulphate  requires  about  400,000  parts  of  water  for  its  solution, 
but  it  is  more  soluble  in  hydrochloric  acid,  even  very  dilute.  In  many  salt 
solutions  it  is  still  more  soluble  than  in  water  acidified  with  hydrochloric  acid. 

2.  In  the  precipitation  of  sulphuric  acid  with  barium  chloride,  the  solu- 
tion should  contain  only  salts  of  the  alkali  metals  and  ammonium,  and  it 
should  be  free  from  nitrates  and  chlorates.     Even  alkali  salts  and  barium 
chloride  are  carried  down  to  some  extent  by  barium  sulphate,  more  or  less 
in  proportion  to  their  concentration,  and  consequently  the  solution  should 
be  dilute.     Since,  further,  the  solubility  of  barium  sulphate,  as  well  as  the 
amount  of  barium  chloride  carried  down,  increases  with  the  concentration 
of  the  hydrochloric  acid  present,  the  quantity  of  the  latter  should  be  re- 
duced to  a  minimum;  some,  however,  must  be  present,  since  otherwise 
the  precipitate  would  be  very  fine  grained,  and  therefore  difficult  to  filter. 

Barium  sulphate  carries  down  quantities  of  chlorine  varying  from  traces 
to  as  much  as  i%,  depending  upon  the  conditions,  and  these  should  there- 
fore be  very  carefully  regulated.  For  a  quantity  of  sulphate  corresponding 
to  1-2  g.  of  BaSC>4,  the  latter  should  be  precipitated  from  a  solution  which 
has  been  diluted  to  about  400  cc.,  and  which  should  contain  1.5  cc.  of  free 
6-normal  hydrochloric  acid.  This  solution  should  be  boiling  hot,  and,  for 
each  gram  of  barium  sulphate,  10  cc.  of  i-normal  barium  chloride  solution 
diluted  to  100  cc.,  and  boiling  hot,  should  be  poured  in,  all  at  once,  with  con- 
stant stirring.  In  this  way  exact  results  can  be  obtained,  but  only  owing  to  a 
compensation  of  errors :  although  a  very  small  quantity  of  barium  sulphate 
remains  dissolved  in  the  acidified  salt  solution,  an  approximately  equal 
weight  of  barium  chloride  is  contained  in  the  ignited  precipitate. 

3.  Owing  to  the  tendency  of  the  precipitate  to  pass  through  the  pores 
of  the  filter,  the  filtrate  and  washings  should  always  be  carefully  examined 
for  minute  quantities  of  the  sulphate.     This  is  best  accomplished  by  im- 
parting to  the  liquid  a  gentle  rotary  motion,  so  that,  if  present,  the  sulphate 
will  collect  at  the  center  of  the  beaker. 

4.  A  partial  reduction  of  barium  sulphate  to  sulphide  may  be  caused 
by  the  action  of  the  burning  carbon  of  the  filter ;  in  order  to  prevent  the 
reduction  of  any  considerable  quantity,  the  crucible  should  not  be  heated 


GRAVIMETRIC  ANALYSIS  65 

above  dull  redness  until  after  the  carbon  has  been  consumed.  Subsequent 
ignition,  with  ready  access  of  air,  will  then  suffice  to  reconvert  the  sulphide 
to  sulphate.  If  considerable  sulphate  is  reduced,  it  may  be  necessary  to 
moisten  the  precipitate  with  sulphuric  acid,  and  then  to  heat  cautiously 
until  the  excess  of  acid  is  expelled. 

THE  DETERMINATION  OF  SULPHUR  IN  ORES 

Method.  The  ore  is  heated  with  strong  nitric  acid  and  potas- 
sium chlorate,  which  oxidize  the  sulphur  to  sulphuric  acid. 
After  removing  the  nitric  acid  and  chlorate,  as  well  as  the  iron, 
lead,  etc.,  the  sulphuric  acid  is  precipitated  with  barium  chloride, 
and  weighed  as  barium  sulphate. 

Procedure.  Treat  0.25-0.50  g.  samples  of  the  finely  pul- 
verized ore  (depending  upon  the  sulphur  content)  in  250  cc. 
Erlenmeyer  flasks  with  10  cc.  of  nitric  acid  (sp.  gr.  142),  and 
heat  very  gently  until  the  red  fumes  have  somewhat  abated. 
Then  increase  the  heat,  and  add  to  the  quietly  boiling  liquid 
potassium  chlorate,  from  time  to  tune,  in  o.i  g.  portions,  until 
any  free  sulphur  which  has  separated  is  entirely  oxidized  and 
dissolved;  finally  add  0.5  g.  of  solid  sodium  chloride  and  evapo- 
rate the  solution  to  dryness.  After  cooling,  cautiously  add 
10  cc.  of  hydrochloric  acid  (sp.  gr.  1.19),  heat  gently  until 
solution  is  as  complete  as  possible,  and  evaporate  to  dryness. 
Take  up  in  5  cc.  of  strong  hydrochloric  acid,  heat  to  boiling, 
and  dilute  with  100  cc.  of  cold  water.  To  the  cold  solution  add 
three  drops  of  methyl  orange,  and  ammonia  to  alkaline  reaction ; 
then  add  5  cc.  more  of  ammonia  and  10  cc.  of  ammonium  car- 
bonate solution.  Heat  to  boiling,  allow  the  precipitate  to  settle 
in  the  hot  liquid,  and  filter  while  still  hot,  washing  thoroughly 
with  hot  water,  and  receiving  the  filtrate  and  washings  in  a  yoo-cc. 
beaker.  Neutralize  the  filtrate  with  hydrochloric  acid,  and  add  i .  5 
cc.  of  the  6-normal  acid  in  excess.  Dilute  the  solution  to  400  cc., 
heat  to  boiling,  and  add  with  stirring  a  boiling-hot  mixture  of  10 
cc.  of  i -normal  barium  chloride  solution  and  100  cc.  of  water. 
Allow  the  mixture  to  stand  for  half  an  hour,  test  for  complete  pre- 


66     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

cipitation,  and  finish  the  determination  as  described  in  the  pre- 
vious Procedure.   Report  the  percentage  of  sulphur  in  the  sample. 

NOTES.  —  i.  Barium  sulphate,  if  present  in  the  ore,  remains  practically 
unaffected  by  the  above  acid  treatment.  If  it  is  desired  to  determine  the 
total  sulphur  hi  ores  containing  barium,  the  hydrochloric  acid  solution, 
after  the  removal  of  nitrates  and  chlorates  and  dilution  with  100  cc.  of 
water,  may  be  treated  with  5  g.  of  solid  ammonium  chloride  (to  hold  any 
lead  in  solution),  heated  to  boiling,  and  filtered  from  the  insoluble  residue. 
The  filter  containing  the  latter  is  destroyed  by  ignition  in  a  platinum  crucible, 
and  the  residue  fused  with  an.excess  of  sodium  carbonate.  The  fusion  is 
extracted  with  hot  water  and  the  residue  washed  with  sodium  carbonate 
solution;  the  filtrate  and  washings,  which  contain  sodium  sulphate,  are 
added  to  the  hydrochloric  acid  filtrate  containing  the  bulk  of  the  sulphur. 
The  united  filtrates  are  then  treated  with  ammonia  and  ammonium  car- 
bonate, as  described  above. 

2.  The  small  amount  of  sodium  chloride  is  added  before  the  first  evapora- 
tion in  order  to  prevent  the  possible  loss  of  any  free  sulphuric  acid  which 
might  be  present,  —  as,  for  example,  in  the  analysis  of  pyrites.    The  potas- 
sium chlorate  added  would,  however,  probably  be  sufficient  in  most  cases 
to  accomplish  this  result.    In  the  analysis  of  pyrites,  which  contains  a 
very  high  percentage  of  sulphur,  samples  should  be  used  of  only  0.25  g. 
Otherwise  the  procedure  is  the  same. 

3.  Upon  adding  ammonia  in  excess  to  the  solution  and  heating  to  boiling, 
there  is  practically  no  danger  of  losing  sulphur  in  the  form  of  basic  ferric 
sulphate.    (In  this  connection  see  Note  3  of  the  foregoing  Procedure.)    The 
ammonium  carbonate  is  added  in  order  to  remove  any  lead  which  may  be 
present,  as  the  carbonate,  and  thus  prevent  the  loss  of  sulphur,  as  PbSC>4, 
before  the  precipitation  with  barium  chloride. 

4.  In  neutralizing  a  solution  with  the  use  of  methyl  orange,  the  solution 
should  be  cold,  since  otherwise  the  methyl  orange  is  not  a  satisfactory 
indicator. 

5.  The  student  should  be  sure  to  read  the  notes  on  the  determination 
of  sulphur  in  iron  sulphate,  and  also  Problems  vi,  12  and  13,  of  Part  IV. 

THE  DETERMINATION  OF  PHOSPHORIC  ANHYDRIDE    IN 
PHOSPHATE  ROCK 

Method.  The  finely  ground  mineral  is  heated  with  nitric 
acid,  the  mixture  evaporated  to  dryness,  and  the  residue  ex- 
tracted with  hot  nitric  acid;  the  solution  is  then  filtered  from 


GRAVIMETRIC  ANALYSIS  67 

the  insoluble  silicious  material.  The  filtrate  is  made  almost 
neutral  with  ammonia,  and  is  treated  with  a  solution  of  am- 
monium molybdate,  in  excess,  to  separate  the  phosphoric  acid 
from  calcium,  iron,  aluminum,  etc.  The  precipitated  ammonium 
phospho-molybdate  is  washed  with  acidified  ammonium  nitrate 
solution,  dissolved  in  ammonium  hydroxide,  and  the  phosphoric 
acid  precipitated  with  magnesia  mixture.  The  magnesium  am- 
monium phosphate  is  finally  ignited  to  magnesium  pyrophos- 
phate,  which  is  weighed. 

Procedure.  Weigh  out  two  portions  of  the  finely  ground 
mineral,  not  to  exceed  0.2  g.  each,  into  2oo-cc.  beakers,  and  treat 
each  as  follows.  Pour  over  the  sample  20  cc.  of  6-normal  nitric 
acid  and  warm  gently  until  solvent  action  has  ceased;  then 
evaporate  the  mixture  to  dryness  on  the  steam  bath.  Allow 
the  dry  residue  to  remain  for  half  an  hour  on  the  steam  bath, 
and  then  heat  it  again  for  a  few  moments  with  20  cc.  of  the 
nitric  acid.  Filter  off  any  siliceous  residue  and  wash  several 
times  with  small  portions  of  hot  water,  receiving  the  filtrate  and 
washings  in  a  4oo-cc.  beaker.  Finally  test  the  washings  with 
ammonia  for  calcium  phosphate,  but  add  to  the  original  filtrate 
all  such  test  solutions  in  which  a  precipitate  appears.  Cautiously, 
and  with  stirring,  add  ammonia  to  the  filtrate  and  washings  until 
the  precipitate  just  fails  to  redissolve,  then  nitric  acid,  drop  by 
drop,  until  the  cloudiness  disappears.  The  volume  of  the  liquid 
at  this  point  should  not  exceed  100  cc. 

Heat  the  solution  until  it  cannot  be  held  comfortably  in  the 
hand,  remove  the  burner,  and  add  75  cc.  of  a  freshly  filtered 
solution  of  ammonium  molybdate  which  has  been  gently  warmed. 
Digest  for  an  hour  at  60-65°,  an<^  then  decant  the  supernatant 
liquid  through  a  filter ;  wash  the  yellow  precipitate  by  decanta- 
tion  with  acid  ammonium  nitrate  solution1  (still  keeping  the 
bulk  of  the  precipitate  in  the  beaker),  until  3  cc.  of  the  washings 
give  no  test  for  calcium  with  ammonia  and  ammonium  oxalate. 

1  Made  by  mixing  50  cc.  of  6-normal  ammonium  hydroxide  with  100  cc.  of 
6-normal  nitric  acid,  and  diluting  the  mixture  with  350  cc.  of  water. 


68     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  filtrate  should  not  be  thrown  away,  but  should  be  tested 
for  complete  precipitation  by  renewed  digestion  with  5  cc.  of 
moybdate  reagent ;  it  should  then  be  placed  in  a  receptacle  for 
"  Molybdate  Residues." 

Dissolve  the  ammonium  phospho-molybdate  by  pouring  over 
the  filter  four  separate  10  cc.  portions  of  a  warm  2.5  per  cent 
solution  of  ammonia  l  (afterwards  washing  the  filter  five  tunes 
with  10  cc.  portions  of  hot  water),  and  receiving  the  filtrate  and 
washings  in  the  beaker  containing  the  bulk  of  the  precipitate. 
To  the  clear  solution  add  hydrochloric  acid,  drop  by  drop,  with 
stirring,  until  the  yellow  cloudiness  produced  disappears  only 
slowly  upon  stirring.  To  this  solution  add  20  cc.  of  magnesia 
mixture  from  a  pipette,  at  the  rate  of  about  i  drop  per  second, 
with  vigorous  stirring  (see  Note  7  ;  the  last  10  cc.  may  be  added 
somewhat  faster).  Let  stand  for  15  minutes,  add  15  cc.  of  am- 
monia (sp.  gr.,  0.90),  and  then,  after  a  period  of  2  or  3  hours,  de- 
cant the  clear  liquid  through  a  filter  and  transfer  the  precipitate 
to  the  filter  by  means  of  2.5  per  cent  ammonia  water.  Continue 
the  washing  with  this  liquid  until  3  cc.  of  the  washings,  after 
acidification  with  nitric  acid,  give  no  opalescence  with  silver 
nitrate  solution.  Finally  test  the  filtrate  for  complete  precipi- 
tation. 

Dry  the  filter  in  the  covered  funnel,  and  then  ignite,  being 
careful  to  raise  the  temperature  slowly  and  to  insure  the  presence 
of  plenty  of  air.  Do  not  raise  the  temperature  above  moderate 
redness  until  the  precipitate  is  white.  Finally  ignite  to  constant 
weight  at  the  blast  lamp,  over  a  large  Meker  burner,  or,  prefer- 
ably, in  an  electric  furnace.  Report  the  percentage  of  P205  in 
the  sample. 

NOTES.  —  i.  The  dehydration  and  removal  of  any  dissolved  silicic  acid 
is  necessary,  since  otherwise  it  would  tend  to  partially  separate  with  the 
phospho-molybdate  precipitate,  and  render  the  latter  more  or  less  insoluble 
in  ammonia. 

1  Made  by  mixing  10  cc.  of  6-normal  ammonium  hydroxide  with  30  cc.  of  hot 
water. 


GRAVIMETRIC  ANALYSIS  69 

When  washing  the  siliceous  residue  the  filtrate  may  be  tested  for  calcium 
by  simply  adding  ammonia,  which  neutralizes  the  acid  holding  calcium  phos- 
phate in  solution  and  causes  precipitation. 

2.  Nitric  acid  is  used  as  the  solvent  because  the  phospho-molybdate  is 
somewhat  soluble  in  hydrochloric  acid.    Nitric  acid  exerts  a  slight  solvent 
action,  but  this  is  counteracted  by  the  presence  of  ammonium  nitrate; 
hence  the  partial  neutralization  of  the  nitric  acid  with  ammonia,  and  the 
washing  with  nitric  acid  containing  an  equivalent  of  ammonium  nitrate. 

It  should  be  noted  that  the  molybdate  reagent  contains  ammonium  nitrate 
and  free  nitric  acid.  (See  the  Preparation  of  Reagents,  in  the  Appendix.) 

3.  The  precipitation  of  the  phosphoric  acid  as  magnesium  ammonium 
phosphate  from  the  original  solution  of  the  rock  is  not  possible,  owing  to 
the  presence  of  metals  such  as  iron,  aluminum,  calcium,  etc.,  which  form 
phosphates  insoluble  in  ammonia.    For  that  reason  the  phosphoric  acid  is 
first  separated  from  the  metals,  in  the  presence  of  nitric  acid,  by  means  of 
ammonium  molybdate. 

4.  While  the  composition  of  the  yellow  precipitate  varies  somewhat 
with  the  conditions,  it  nevertheless  seems  to  correspond  pretty  closely  to 
the  formula  (NH4)3P04 . 12  Mo03 .  2  HNO3 .  H20 ;  at  any  rate,  the  ratio 
P2O5 :  24  Mo03  holds  good.    The  yellow  precipitate  dissolves  in  ammonia 
to  give  ammonium  phosphate  and  ammonium  molybdate,  and  molybdic 
acid  is  not  precipitated  by  magnesia  mixture. 

5.  The  precipitation  of  the  phospho-molybdate  is  more  prompt  in  warm 
than  in  cold  solutions,  but  the  temperature  should  not  exceed  65° ;  at  higher 
temperatures  molybdic  acid,  which  is  white,  tends  to  separate.    Vigorous 
stirring  also  promotes  the  separation  of  the  yellow  precipitate. 

A  large  excess  of  the  molybdate  reagent  is  required  to  effect  a  complete 
precipitation  of  the  phosphoric  acid.  Theoretically  1.95  g.  of  Mo03  are 
required  to  combine  with  the  phosphorus  in  0.2  g.  of  rock  containing  40 
per  cent  of  P2O5 ;  while  the  quantity  of  the  reagent  actually  used  (75  cc.) 
contains  about  5  g.  of  MoOg.  The  presence  of  ammonium  nitrate  in  the 
solution  is  also  conducive  to  complete  precipitation.  These  substances, 
by  mass  action,  prevent  the  partial  dissociation  of  the  complex  into  its  more 
soluble  constituents. 

6.  If  magnesium  ammonium  phosphate  is  washed  with  pure  water,  it  is 
hydrolyzed  according  to  the  equation, 

MgNH4P04+H20  ^  MgHPO4+NH4OH. 

It  thereby  loses  its  crystalline  form,  and  is  almost  sure  to  give  a  cloudy 
filtrate.  In  the  presence  of  ammonium  hydroxide,  however,  this  decom- 
position is  prevented  by  mass  action. 


70     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

7.  The  precipitate  of  magnesium  ammonium  phosphate  should  be  per- 
fectly crystalline ;  the  slow  addition  of  the  reagent,  with  constant  stirring, 
is  essential  to  this  end,  but  the  stirring  rod  should  not  be  allowed  to  scratch 
the  beaker. 

A  large  excess  of  magnesia  mixture  tends  to  cause  the  precipitate  to  carry 
down  molybdic  acid,  as  well  as  magnesia  (shown  by  a  persistently  flocculent 
precipitate).  In  such  cases  the  precipitate  should  be  redissolved  by  adding 
a  small  quantity  of  hydrochloric  acid,  the  solution  treated  with  2  cc.  of 
magnesia  mixture,  and  the  hot  liquid  slowly  neutralized  with  2.5  per  cent 
ammonia.  Strong  ammonia  is  then  added,  and  the  analysis  continued  as 
above. 

"Magnesia  Mixture"  is  prepared  by  putting  together  in  solution  mag- 
nesium chloride,  ammonium  chloride,  and  ammonium  hydroxide.  The 
function  of  the  ammonium  chloride  is  to  prevent  the  precipitation  of  mag- 
nesium hydroxide,  so  that  the  composition  of  the  precipitate  may  corre- 
spond to  the  formula,  MgNH4P04 .  6  H20. 

8.  Upon  ignition,  the  magnesium  ammonium  phosphate  gives  off  am- 
monia  and   water,    and   is    converted    into    magnesium    pyrophosphate: 
2NH4MgP04.6H20  =  Mg2P207+2NH3+i3H2O.    The   precautions   de- 
tailed hi  Part  I  should  be  observed  with  great  care  during  the  ignition  of 
this  precipitate.    There  is  danger  here  of  a  partial  reduction  of  the  phos- 
phate by  the  ammonia  or  by  the  carbon  of  the  filter,  and  also,  if  too  soon 
heated  very  strongly,  the  precipitate  becomes  glazed  over  and  it  is  then 
practically  impossible  to  remove  the  carbon  by  further  heating.     The 
precipitate  is  much  more  readily  ignited  to  whiteness  in  platinum  than 
hi  porcelain ;  but  in  case  platinum  is  used,  especial  care  should  be  taken  to 
provide  a  plentiful  supply  of  air.    Reduction  of  the  phosphorus  would 
play  havoc  with  the  crucible. 

The  most  satisfactory  procedure  is  to  filter  off  the  magnesium  ammonium 
phosphate  in  a  Munroe  crucible  of  platinum  (with  a  platinum  sponge  filter) , 
in  which  the  precipitate  can  be  ignited  without  danger  of  loss.  If  a  good 
muffle  furnace  (preferably  electric)  is  available,  a  Gooch  crucible  of  porce- 
lain may  be  used  with  advantage. 

THE  DETERMINATION  OF  CALCIUM  AND  MAGNESIUM 
OXIDES  IN  LIMESTONE 

Method.  The  hydrochloric  acid  extract  of  the  limestone x  is 
freed  from  dissolved  silica,  treated  with  bromine  water  and 
ammonia,  to  remove  iron,  aluminum,  manganese,  etc.,  and  from 

1  See  Note  i. 


GRAVIMETRIC  ANALYSIS  71 

the  filtrate  the  calcium  is  precipitated  with  ammonium  oxalate, 
the  precipitate  being  ignited  to  the  oxide.  The  filtrate  from 
the  calcium  oxalate  is  treated  with  sodium  phosphate  and  am- 
monia, the  precipitate  of  magnesium  ammonium  phosphate 
being  ignited  to  magnesium  pyro-phosphate. 

A .  Procedure  for  the  Determination  of  Calcium.  Weigh  out 
into  two  casseroles  0.5-0.6  g.  samples  of  the  finely  ground  rock, 
and  treat  each  as  follows :  Cautiously  moisten  the  powder  with 
5  cc.  of  water,  cover  the  casserole,  add  10  cc.  of  6-normal 
hydrochloric  acid  in  small  portions,  and  evaporate  to  dryness 
on  the  steam  bath.  Pour  over  the  residue  5  cc.  of  water  and 
10  cc.  of  the  hydrochloric  acid,  evaporate  to  dryness,  and  heat 
the  dry  residue  for  half  an  hour  on  the  steam  bath.  Pour  over 
this  residue  5  cc.  of  wate'r  and  10  cc.  of  the  6-normal  acid,  and 
heat  gently  for  10  minutes;  filter  and  wash  twice  with  5  cc. 
portions  of  dilute  hydrochloric  acid,  and  finally  with  water  until 
free  from  chlorides.  (See  Note  i  in  regard  to  the  insoluble  residue.) 

Add  to  the  filtrate  and  washings  enough  bromine  water  to 
impart  a  distinctly  yellow  tinge,  boil,  and  then  add  ammonia 
until  its  odor  persists  in  the  solution.  Heat  until  the  excess  of 
ammonia  is  largely  expelled,  and  filter  promptly.  Wash  the 
filter  twice  with  hot  water,  allowing  the  washings  to  run  into  the 
beaker  containing  the  filtrate.  Now  pour  through  the  filter 
25  cc.  of  hot  hydrochloric  acid  (one  volume  of  the  6^normal 
acid  to  5  of  water),  and  if  there  is  a  brown  insoluble  residue 
allow  it  to  remain  on  the  filter ;  the  acid  solution  should  be  re- 
ceived in  the  beaker  in  which  the  ammonium  hydroxide  pre- 
cipitate was  obtained.  Wash  the  filter  five  times  with  hot 
water,  and  then  reprecipitate  the  iron,  etc.,  from  the  filtrate  and 
washings  with  bromine  water  and  ammonia  as  already  described. 
Collect  the  precipitate  on  the  filter  already  used,  and  wash  it  free 
from  chlorides  with  hot  water ;  add  the  filtrate  and  washings  to 
those  at  first  obtained.  (Concerning  the  precipitate,  see  Note  i.) 

Evaporate  the  combined  filtrates  and  washings  to  a  volume  of 
about  200  cc.  Heat  the  solution  to  boiling;  if  necessary,  add 


72     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

ammonia  until  its  odor  is  plainly  perceptible;  and  then  add 
ammonium  oxalate  solution  slowly  and  with  stirring,  in  moderate 
excess.  Boil  for  two  minutes,  allow  the  precipitate  to  settle 
for  half  an  hour,  and  decant  through  a  filter  into  a  beaker,  wash- 
ing the  precipitate  twice  with  hot  water  containing  a  few  cubic 
centimeters  of  ammonium  oxalate  solution  and  a  very  little 
ammonia.  Test  the  filtrate  with  ammonium  oxalate  for  com- 
plete precipitation,  and  if  no  precipitate  forms  in  15  minutes, 
acidify  the  solution  with  hydrochloric  acid  and  reserve  it  for 
the  magnesium  determination. 

Redissolve  the  calcium  oxalate  with  warm  hydrochloric  acid 
(one  volume  of  the  6-normal  acid  to  one  of  water),  pouring  the 
acid  through  the  filter  and  receiving  it  in  the  beaker  containing 
the  bulk  of  the]  precipitate.  Wash  the  filter  three  times  with 
water,  and  twice  with  very  dilute  ammonia.  Dilute  the  solu- 
tion to  250  cc.,  heat  to  boiling,  add  i  cc.  of  ammonium  oxalate 
solution,  and  ammonia  in  slight  excess;  boil  for  two  minutes, 
and  set  aside  for  half  an  hour.  Filter  off  the  precipitate  upon 
the  filter  previously  used,  and  wash  it  free  from  chlorides  with 
hot  water  containing  a  few  drops  of  ammonium  oxalate  solution 
and  a  very  little  ammonia.  (The  filtrate  and  washings  should 
at  once  be  acidified  with  hydrochloric  acid  and  combined  with 
those  from  the  first  precipitation.  If  not  already  started,  the 
evaporation  of  these  filtrates  should  be  begun  at  this  point.) 

Gently  ignite  the  dried  precipitate  and  filter  until  the  latter 
is  consumed,  and  then  heat  with  the  full  flame  of  the  burner  for 
45  minutes ;  finally  heat  for  three  minutes  over  the  blast  lamp. 
Repeat  the  heating  at  the  blast  lamp,  until  the  weight  becomes 
constant.  (A  muffle  furnace  is  preferable  for  the  ignition,  after 
the  filter  is  consumed.)  Report  the  percentage  of  CaO  found. 

NOTES.  —  i.  The  chief  component  of  limestone,  calcium  carbonate,  is 
readily  attacked  by  hydrochloric  acid,  as  also  are  some  of  the  other  com- 
ponents ;  but  few  limestones  are  so  pure  as  to  dissolve  completely  in  hydro- 
chloric acid.  The  residue  may  contain  quartz,  silicates,  pyrites,  or  other 
refractory  materials,  and  carbonaceous  matter  may  also  be  present. 


GRAVIMETRIC  ANALYSIS  73 

Furthermore,  the  insoluble  silicates  of  the  residue  are  apt  to  contain  some 
calcium  and  magnesium.  The  thorough  analysis  of  a  limestone  necessi- 
tates the  use  of  an  elaborate  system  of  procedures  and  the  determination 
of  numerous  substances,  but  for  many -  technical  purposes  the  analysis 
may  be  confined  to  the  determination  of  the  insoluble  matter  and  silica, 
of  the  oxides  of  iron  and  aluminum  (including  small  quantities  of  the  oxides 
of  titanium,  manganese,  phosphorus,  etc.),  of  calcium  oxide,  and  of  mag- 
nesium oxide. 

In  this  exercise,  the  insoluble  residue,  if  ignited  and  weighed,  would  give 
a  more  or  less  accurate  approximation  of  the  insoluble  matter  and  silica, 
and  the  ignited  ammonium  hydroxide  precipitate  would  roughly  approxi- 
mate the  summation  of  the  oxides  of  iron,  aluminum,  titanium,  etc.,  in  the 
soluble  portion.  It  should  be  remembered,  however,  that  substances  such 
as  hydrous  silicates,  pyrites,  and  carbonaceous  matter  in  the  insoluble 
residue,  and  ferrous  and  manganous  oxides  (originally  present  as  carbonates) 
in  the  soluble  portion,  would  not  be  correctly  indicated  by  this  method. 

It  should  be  noted  that  the  amount  of  insoluble  residue  and  also  its 
character  will  often  depend  upon  the  concentration  of  the  acid  used  for  the 
solution  of  the  limestone,  and  that  the  determination  of  this  residue  is 
essentially  empirical. 

For  a  description  of  the  complete  analysis  of  limestones,  the  student 
should  refer  to  Bulletin  422  of  the  United  States  Geological  Survey,  by 
W.  F.  Hillebrand. 

2.  Some  of  the  silicates  present  are  apt  to  be  at  least  partly  decomposed 
by  the  acid,  and  the  soluble  silicic  acid  must  be  dehydrated  and  rendered 
insoluble  by  evaporation  and  heating.    The  residue  is  washed  first  with 
dilute  acid  to  prevent  the  separation  on  the  filter  of  basic  salts  of  iron, 
aluminum,  etc.,  owing  to  the  hydrolytic  action  of  water. 

3.  The  addition  of  bromine  water  serves  to  oxidize  any  ferrous  iron, 
and  also  manganese,  which  precipitates  as  MnO(OH)2.    The  ammonium 
hydroxide  precipitate  should  be  filtered  off  promptly,  since  the  alkaline 
solution  absorbs  carbon  dioxide  from  the  air,  with  the  consequent  pre- 
cipitation of  a  little  calcium  as  the  carbonate.    This  is  always  possible, 
and  for  that  reason,  as  well  as  because  the  precipitated  hydroxides  also 
tend  to  carry  down  the  hydroxides  of  calcium  and  magnesium,  the  precipi- 
tate is  redissolved  and  again  precipitated,  to  free  it  from  these  metals. 

4.  The  accurate  separation  of  calcium  and  magnesium  by  means  of 
ammonium  oxalate  requires  considerable  care.    The  calcium  oxalate  tends 
to  carry  down  some  magnesium  oxalate,  probably  in  the  form  of  a  double 
salt,  but  this  can  be  removed  by  dissolving  the  precipitate  and  reprecipi- 
tating  the  calcium  in  the  presence  of  only  this  small  amount  of  magnesium. 


74     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

If  the  proportion  of  magnesium  is  not  very  large,  the  calcium  can  be  sepa- 
rated by  precipitation  from  a  rather  dilute  solution,  with  the  addition  of 
more  than  enough  ammonium  oxalate  to  convert  both  the  magnesium  and 
calcium  into  oxalates.  (In  this  connection,  see  T.  W.  Richards,  C.  T. 
McCaffrey,  and  H.  Bisbee;  Zeitschrift  fiir  anorganische  Chemie,  28,  p.  71 
(1901).) 

5.  The  small  quantity  of  ammonium  oxalate  solution  is  added  before 
the  second  precipitation  of  the  calcium,  because  an  excess  of  the  reagent 
reduces  the  solubility  of  the  calcium  oxalate,  and  also  tends  to  hold  the 
magnesium  in  solution  in  the  form  of  a  double  magnesium  ammonium 
oxalate.    For  the  first  reason,  the  precipitate  is  washed,  not  with  pure  water, 
but  with  water  containing  ammonium  oxalate  and  ammonia.    These  sub- 
stances are  volatilized  in  the  ignition. 

6.  Calcium  oxalate  is  practically  insoluble  in  water  (5.6  mg.  of  the 
anhydrous  salt  per  liter  of  saturated  solution),  and  only  very  slightly  soluble 
in  acetic  acid,  but  it  is  readily  dissolved  by  the  strong  mineral  acid?.    This 
behavior  with  acids  is  explained  by  the  fact  that  oxalic  acid  lies  about 
halfway  in  strength  between  acetic  acid  and  the  strong  mineral  acids.     In 
acetic  acid  solution,  the  hydrogen-ion  concentration  is  too  low  to  appre- 
ciably diminish  the  concentration  of  C204™  ions,  and  practically  no  solvent 
action  takes  place.     In  the  solution  of  a  strong  mineral  acid,  however,  the 
high  hydrogen-ion  concentration  causes  the  calcium  oxalate  to  dissolve 
according  to  the  following  scheme : 

CaC2O4  ^  CaC204  ^  Ca+++  C2O4~  1         f  HC204-  or 
(solid)  (dissolved)  |  ^T 

HC1    $  Cl-  +    H+     J         I H2C204. 

The  oxalate  is  immediately  reprecipitated  from  such  a  solution  upon  the 
addition  of  a  base;  the  hydroxide  ions  unite  with  the  hydrogen  ions  of 
both  the  mineral  acid  and  the  oxalic  acid  to  form  water,  and  the  Ca++  and 
C204 —  ions  left  in  the  solution  recombine  to  form  CaC204.  (Compare 
the  precipitation  of  Ca3(P04)2  from  the  acid  solution  of  apatite  by  am- 
monium hydroxide,  in  the  preceding  exercise.) 

7.  Upon  ignition,  calcium  oxalate  becomes  anhydrous  slightly  above 
1 80°,  and  at  low  redness  it  is  decomposed  into  calcium  carbonate  and 
carbon  monoxide.     Strong  ignition  converts  the  carbonate  into  the  oxide ; 
in  a  platinum  crucible,  this  conversion  may  be  carried  to  completion  over 
a  large  Meker  burner.    With  porcelain  crucibles,  however,  an  electric  fur- 
nace is  to  be  preferred. 

Since  calcium  oxide  absorbs  moisture  and  carbon  dioxide  from  the  air, 
it  should  be  weighed  as  rapidly  as  possible. 


GRAVIMETRIC  ANALYSIS  75 

8.  By  burning  off  the  filter  and  then  evaporating  with  2-3  cc.  of  6-normal 
sulphuric  acid,  the  calcium  oxalate  may  be  converted  into  sulphate,  and 
this  may  be  heated  to  constant  weight.  While  this  procedure  is  preferred 
by  some  analysts,  it  is  nevertheless  disadvantageous,  since  it  involves  danger 
of  loss  by  spattering.  Moreover,  calcium  sulphate  is  more  readily  decom- 
posed upon  ignition  than  barium  sulphate,  and  there  is  some  danger  of  loss 
on  this  account. 

B.  Procedure  for  the  Determination  of  Magnesium.  Evapo- 
rate the  acidified  filtrates  and  washings  from  the  calcium  oxalate 
on  the  steam  bath  until  the  salts  begin  to  crystallize.  Dilute 
the  solution  cautiously  with  small  portions  of  water,  and  with 
stirring,  until  the  salts  are  brought  back  into  solution,  adding 
a  little  hydrochloric  acid  if  the  solution  has  been  evaporated  to 
a  very  small  volume.  (If  the  acid  solution  contains  solid  matter 
at  this  point,  it  should  be  filtered.)  Carefully  add  ammonia 
to  the  clear  solution,  just  to  alkaline  reaction  (methyl  orange) ; 
add  sodium  ammonium  phosphate  solution,  drop  by  drop  with 
stirring,  as  long  as  a  precipitate  continues  to  form,  and  then  10  cc. 
in  excess.  Finally  add  to  the  solution  one  third  of  its  volume  of 
6-normal  ammonia,  stir  vigorously  for  10  minutes,  and  allow  the 
mixture  to  stand  for  at  least  6  hours,  —  preferably  overnight. 

Decant  the  solution  through  a  filter,  and  wash  the  bulk  of 
the  precipitate  on  to  the  filter  with  2.5  per  cent  ammonia  (one 
volume  of  6-normal  ammonia  to  three  of  water) ;  do  not  bother 
to  clean  the  beaker  completely.  Dissolve  the  precipitate  from 
the  filter  with  the  least  possible  quantity  of  hydrochloric  acid 
(6-normal  acid  diluted  with  twice  its  volume  of  water),  receiving 
the  acid  solution  in  the  precipitation  beaker.  Wash  the  filter 
with  small  portions  of  hot  water  until  the  washings  are  free  from 
chlorides.  Add  to  the  combined  filtrate  and  washings  2  cc.  of 
sodium  ammonium  phosphate  solution  and  then  aqueous  am- 
monia, drop  by  drop  with  constant  stirring,  until  the  liquid 
smells  distinctly  of  ammonia.  Stir  for  two  minutes,  add  to  the 
solution  one  third  its  volume  of  6-normal  ammonia,  and  allow 
the  mixture  to  stand  for  2  hours.  Decant  the  clear  liquid  through 
a  filter  and  transfer  the  precipitate  to  the  filter  by  means  of  2.5  per 


76     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

cent  ammonia.  Continue  the  washing  with  this  liquid  until  3  cc.  of 
the  washings  give  no  opalescence  with  nitric  acid  and  silver  nitrate. 
Dry  the  filter  completely  in  the  covered  funnel,  and  then 
ignite,  taking  great  pains  to  raise  the  temperature  very  slowly 
and  to  insure  the  presence  of  plenty  of  air.  Do  not  raise  the 
temperature  above  moderate  redness  until  the  precipitate  is 
white.  Finally  ignite  to  constant  weight  over  the  blast  lamp 
or  in  a  muffle  furnace.  Report  the  percentage  of  MgO  found. 

NOTES.  —  i.  The  filtrates  from  the  calcium  oxalate  should  be  slightly 
acidified  immediately  after  filtration,  in  order  to  avoid  the  solvent  action 
of  the  alkaline  solution  upon  glass. 

2.  The  precipitation  of  the  magnesium  should  be  made  in  a  small  volume 
of  liquid,  and  the  ratio  of  ammonia  to  the  total  volume  of  solution  should 
be  carefully  regulated,  on  account  of  the  relative  solubility  of  the  magnesium 
ammonium  phosphate.    (Compare  Note  6,  under  the  determination  of  Phos- 
phoric Anhydride.) 

3.  In  the  presence  of  ammonium  salts  in  large  quantity,  the  first  precipi- 
tate is  rarely  wholly  crystalline ;  it  is  apt  to  contain  the  mixed  phosphate, 
Mg[(NH4)2PO4]2,  which  upon  ignition   leaves   magnesium   metaphosphate 
Mg(P03)2;  and,  if  produced  in  the  presence  of  too  much  ammonia,  it  may 
also  contain  some  tri-magnesium  phosphate,  which  upon  ignition  remains 
unchanged  as  Mg3(PO4)2.     Such  precipitates  can  be  purified  by  dissolving 
in  a  very  little  hydrochloric  acid,  adding  a  small  quantity  of  sodium   am- 
monium phosphate,  and  reprecipitating  the  magnesium,  in  the  practical 
absence  of  ammonium  salts,  by  means  of  ammonia. 

4.  In  order  to  avoid  a  partial  reduction  of  the  phosphorus,  the  precipi- 
tate should  be  heated  gently  at  first,  until  all  the  ammonia  is  expelled  and 
the  filter  consumed.     Also,  if  heated  too  soon  to  a  bright  red  heat,  the 
precipitate  becomes  glazed  over,  and  it  is  then  impossible  to  remove  the 
carbon  by  further  heating. 

Concerning  the  use  of  Munroe  or  of  Gooch  crucibles,  see  Note  8  under 
the  determination  of  phosphoric  anhydride. 

THE  DETERMINATION  OF  CARBON  DIOXIDE  IN 
LIMESTONE 

Method.  The  weighed  carbonate  is  placed  in  an  apparatus 
which  contains  acid  in  a  separate  compartment;  the  whole 
apparatus  is  then  weighed.  After  this  the  acid  is  run  in  upon 


GRAVIMETRIC  ANALYSIS  77 

the  carbonate,  and  the  carbon  dioxide  set  free  is  removed  from 
the  apparatus  through  a  tube  filled  with  calcium  chloride,  which 
prevents  the  escape  of  moisture  from  the  apparatus.  Finally, 
the  apparatus  is  weighed  again,  and  the  loss  in  weight  indicates 
the  quantity  of  carbon  dioxide  in  the  sample. 

Many  different  forms  of  apparatus  have  been  devised  for  this 
purpose.  The  one  shown  in  the  accompanying  figure  is  an  im- 
proved form  of  the  so-called  alkalimeter  of  Mohr.  It  consists 
of  a  small,  wide-mouthed,  flat-bottomed  flask  F,  which  has  a 
ground-glass  connection  with  the  tubes  A  and  B,  which  are  for 
acid  and  calcium  chloride.  The  ground-glass  joints  are  lubri- 
cated with  a  mixture  of  vaseline  and  beeswax,  or  other  suitable 
substance. 

Procedure.  Thoroughly  clean  the  apparatus,  allow  it  to  drain, 
and  finally  dry  it  by  gently  heating  the  flask  while  drawing  a 
current  of  dry  air  through  it.  As  aspirator  an  inverted  wash- 
bottle  (shown  in  the  figure  on  a  very  much  reduced  scale)  may  be 
used,  from  which  the  water  is  caused  to  run  out  slowly  through 
the  shorter  tube.  During  aspiration  the  calcium  chloride  tubes 
C  and  D  should  be  connected  with  c  and  d,  as  shown  in  the 
figure,  so  that  no  moisture  may  enter  the  apparatus.  After 
drying  the  apparatus,  place  a  loose  wad  of  cotton  at  the  bottom 
of  B ;  introduce  into  the  neck  of  the  tube  a  cylinder  of  glazed 
paper  about  3  cm.  wide,  and  through  this  cylinder  pour  in  small 
pieces  of  calcium  chloride  until  the  tube  is  about  three  fourths 
full;  remove  the  glazed  paper,  taking  care  to  keep  the  upper 
walls  of  the  tube  free  from  calcium  chloride.  Place  another 
cotton  wad  in  the  tube,  insert  the  stopper,  and  close  the  tube 
temporarily  at  d  by  means  of  a  short  piece  of  glass  rod  within 
rubber  tubing.1 

1  The  tube  must  be  kept  closed  when  not  in  use,  to  prevent  the  gradual  absorp- 
tion of  moisture  from  the  air.  Each  of  the  ordinary  calcium  chloride  tubes  pre- 
viously mentioned  is  filled  in  the  same  way  about  two  thirds  full,  but  in  this  case  a 
softened  cork  stopper,  pierced  by  a  short  piece  of  glass  tubing  with  rounded  ends, 
is  introduced  and  shoved  far  into  the  tube  with  the  help  of  a  stirring  rod,  leaving 
the  outer  2  or  3  mm.  empty.  This  space  in  the  tube  is  filled  with  molten  sealing- 


78     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 


GRAVIMETRIC  ANALYSIS  79 

When  all  is  ready,  weigh  out  into  the  flask  about  1.5  g.  of  the 
finely  powdered  substance,  which  has  been  dried  at  100°  C.  and 
allowed  to  cool  in  a  desiccator,  and  add  3-4  cc.  of  water.  Close 
the  stopcock  T  and  fill  the  tube  A  about  three  fourths  full 
with  hydrochloric  acid  (i  volume  of  6-normal  acid  to  1.5  volumes 
of  water)  by  means  of  a  small  funnel.  The  whole  apparatus, 
with  the  tubes  open  at  c  and  J,  is  now  accurately  weighed ;  the 
two  calcium  chloride  tubes  are  connected  at  c  and  d\  and  the 
stopcock  T  is  slightly  opened  so  that  the  acid  from  A  slowly 
drops  into  the  flask.  As  soon  as  the  evolution  of  carbon  dioxide 
begins  to  take  place  quietly,  the  apparatus  is  allowed  to  stand 
without  watching  for  about  half  an  hour.  All  of  the  acid  will 
then  have  entered  the  flask,  and  the  decomposition  will  be  practi- 
cally complete.  It  now  remains  to  remove  the  carbon  dioxide 
absorbed  by  the  liquid  and  contained  in  the  apparatus.  To 
this  end,  connect  the  calcium  chloride  tube  D  with  the  wash 
bottle  W,  as  shown  in  the  figure  (the  wash  bottle  is  of  course 
much  larger  than  the  figure  would  indicate),  and  regulate  the 
flow  of  water  through  e  so  that  not  more  than  3  or  4  bubbles 
of  air  per  second  pass  through  the  flask  F.  Then  heat  the  flask 
F  gently,  by  means  of  a  small  flame,  until  the  acid  just  begins 
to  boil ;  at  once  remove  the  flame,  and  continue  to  aspirate  air 
through  the  apparatus  until  it  is.  cold.  Stopper  the  tubes  at  c 
and  dj  wipe  the  apparatus  with  a  clean  dry  towel,  and  allow  it 
to  stand  for  one  half  hour  near  the  balance.  Finally,  remove 
the  stoppers  from  c  and  d,  and  weigh  the  apparatus.  Report 
the  percentage  of  C02  found. 

NOTES.  —  i.  This  method  yields  excellent  results  in  the  estimation  of 
large  amounts  of  carbonic  acid  such  as  are  present  in  limestones  and  baking 
powders.  But  it  is  unreliable  for  the  determination  of  small  quantities, 
e.g.  in  cements. 

2.  Since  baking  powders  are  decomposed  by  water,  they  should  be  kept 
dry  until  after  the  apparatus  has  been  weighed ;  and  since  their  efficiency 

wax,  so  that  an  air-tight  connection  is  made.    These  tubes  also  are  closed,  when 
not  in  use,  by  glass  rods  within  rubber  tubing. 


8o    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

as  leavening  agents  depends  upon  the  volume  of  gas  liberated  under  the 
conditions  of  actual  usage,  water  should  be  placed  in  the  tube  A,  instead 
of  acid.  Otherwise  the  procedure  is  the  same.  The  loss  in  weight  is  then 
a  measure  of  the  available  carbon  dioxide  of  the  sample. 

3.  Carbon  dioxide  is  readily  displaced  from  the  apparatus  by  the  method 
described,  but  in  order  to  insure  its  complete  removal  at  least  a  liter  of  air 
should  be  drawn  through  the  apparatus.    The  small  tube  from  A  should 
project  well  below  the  surface  of  the  liquid,  for  in  order  to  remove  carbon 
dioxide  efficiently  from  the  solution,  the  air  must  be  made  to  bubble  through 
the  liquid. 

4.  Since  commercial  calcium  chloride  is  apt  to  contain  free  lime,  it 
should  for  the  best  results  be  treated  with  carbon  dioxide  before  the  deter- 
mination is  made.    For  this  purpose  a  current  of  the  dry  gas  is  passed 
through  the  apparatus  for  a  minute  or  two,  the  tubes  at  c  and  d  are  closed, 
and  the  apparatus  allowed  to  stand  overnight.      The  carbon  dioxide  is 
then  removed  by  aspirating  dry  air  through  the  apparatus  for  about  20 
minutes,  after  which  the  sample  may  be  placed  in  the  flask. 

5.  The  most  serious  objection  to  this  method  is  the  fact  that,  owing  to 
the  size  and  weight  of  the  apparatus,  there  is  likely  to  be  an  appreciable 
error  in  the  difference  between  the  two  weights.    This  danger,  however, 
can  be  largely  overcome  if  a  similar  piece  of  apparatus  is  available  as  a  tare. 

THE  DETERMINATION  OF  SILICA  IN  A  REFRACTORY 

SILICATE 

Method.  The  finely  ground  sample  is  fused  with  an  excess 
of  sodium  carbonate,  whereby  it  yields  sodium  silicate  and  other 
compounds,  depending  upon  the  nature  of  the  mineral.  The 
melt  is  then  decomposed  with  hydrochloric  acid,  which  should 
dissolve  everything  except  a  portion  of  the  silicic  acid.  Upon 
evaporating  the  liquid,  the  silicic  acid  in  solution  loses  water 
and  becomes  much  less  soluble;  upon  extracting  the  residue 
with  hydrochloric  acid,  filtering  from  silica,  and  evaporating 
the  filtrate,  however,  appreciable  amounts  of  silica  are  recovered 
in  a  second  filtration,  leaving  negligibly  small  amounts  in  the 
filtrate.  Upon  strongly  igniting  the  precipitates,  the  silica 
(contaminated  with  iron  oxide  and  alumina)  is  left  in  the 
anhydrous  condition.  After  weighing,  this  is  evaporated  with 
hydrofluoric  acid  and  a  few  drops  of  sulphuric  acid,  and  the 


GRAVIMETRIC  ANALYSIS  81 

residue  is  subjected  to  strong  ignition.  The  weight  of  the  im- 
pure silica  less  that  of  the  ignited  residue  gives  the  weight  of  the 
silica  originally  in  the  sample. 

Procedure.  Grind  about  3  g.  of  the  material  in  an  agate 
mortar  until  it  will  entirely  pass  through  a  sieve  of  fine  silk  bolt- 
ing cloth.1  Weigh  out  into  two  platinum  or  palau  (also  called 
rhotanium)  crucibles  portions  of  the  silicate  of  about  0.75  g. 
each.  Also  weigh  out,  on  a  rough  balance,  two  portions  of 
anhydrous  sodium  carbonate  of  about  4  g.  each.  In  each  case, 
add  about  three  fourths  of  the  sodium  carbonate  to  the  silicate 
sample  in  the  crucible,  place  the  latter  on  a  piece  of  glazed  paper, 
and  thoroughly  mix  its  contents  with  a  dry  glass  rod.  Place  the 
remaining  fourth  of  the  flux  on  top  of  the  mixture,  after  first 
stirring  it  with  the  rod  to  remove  from  the  latter  any  adhering 
particles  of  the  mixture.  Cover  the  crucible  and  heat  it  gradually 
to  the  highest  heat  of  the  Bunsen  or  Tirrill  burner,  and  then,  if 
necessary  to  secure  complete  fusion,  heat  the  mixture  over  a 
Meker  burner  or  a  blast  lamp.  As  soon  as  the  mass  is  in  quiet 
fusion,  evolving  no  gas  bubbles,  take  up  the  crucible  in  tongs 
applied  to  the  upper  edge,  and,  by  means  of  a  slow  rotary  motion, 
cause  the  liquid  melt  to  spread  around  the  walls  of  the  crucible, 
where  it  will  solidify.  As  soon  as  this  takes  place,  and  while 
the  mass  is  still  red-hot,  plunge  the  lower  portion  of  the  crucible 
for  a  few  seconds  into  cold  water,  but  with  care  not  to  allow  any 
water  to  enter  the  crucible.  Then  set  the  crucible  aside  to  cool. 
The  solid  material  may  then  be  loosened  from  the  crucible  by 
gentle  tapping.  (Do  not  deform  the  crucible?) 

Place  the  solid  melt  in  a  rather  tall  beaker,  add  100  cc.  of 
water  and,  with  stirring,  gradually  add  50  cc.  of  6-normal  hydro- 
chloric acid.  Also  clean  the  crucible  and  lid  with  a  little  of  the 
acid,  and  add  this  to  the  main  portion  in  the  beaker.  (In  case 

1  Place  the  ground  material  in  a  small  beaker  and  stretch  over  the  top  a  piece  of 
the  bolting  cloth,  fastening  the  cloth  in  place  by  means  of  a  rubber  band  below 
the  rim  of  the  beaker.  By  gently  tapping  the  inverted  beaker  over  a  piece  of  clean 
paper,  the  fine  particles  are  caused  to  pass  through  the  sieve.  The  coarser  par- 
ticles which  fail  to  pass  through  must  be  returned  to  the  mortar  and  reground. 

G 


82     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  melt  adheres  obstinately  to  the  crucible,  place  both  in  the 
beaker  and  treat  with  water  and  acid  as  described.)  Heat  the 
beaker  gently,  and,  if  necessary,  aid  the  disintegration  of  the 
melt  by  gentle  pressure  with  the  broadened  end  of  a  glass  rod. 
After  complete  disintegration,  transfer  the  mixture  to  a  porce- 
lain casserole,  evaporate  to  dryness  on  the  steam  bath,  stirring 
frequently  towards  the  end  until  the  residue  is  a  dry  powder. 
Heat  the  dry  residue  on  the  steam  bath  for  at  least  an  hour, 
cover  it  with  5  cc.  of  i2-normal  hydrochloric  acid,  warm  gently, 
see  that  the  residue  is  wholly  moistened  with  acid,  add  100  cc. 
of  water,  and  heat  to  boiling.  Filter  promptly,  and  wash  five 
times  with  hot  dilute  acid  (i  volume  of  6-normal  hydrochloric 
acid  to  3  of  water),  collecting  the  nitrate  and  washings  in  a  por- 
celain casserole ;  evaporate  to  dryness,  heat  the  dry  residue  on 
the  steam  bath  for  one  hour,  and  proceed  as  before,  using  a  fresh 
filter  to  remove  the  silica.  Wash  the  filter  with  hot  dilute 
hydrochloric  acid,  as  before,  and  then  wash  both  filters  with 
hot  water  until  the  washings  are  free  from  chlorides. 

Transfer  both  filters  to  a  weighed  platinum  (or  palau)  crucible, 
and  ignite  cautiously  until  the  paper  is  consumed,  then  at  the 
full  heat  of  the  burner  for  half  an  hour.,  Moisten  the  cold 
residue  with  2  or  3  drops  of  strong  sulphuric  acid,  heat  cautiously 
to  expel  the  free  acid,  ignite  to  low  redness,  and  finally  for  half 
an  hour  over  the  blast  lamp.  Repeat  the  blasting  for  periods 
of  5  minutes,  to  constant  weight. 

Now  add  to  the  silica  in  the  crucible  i  cc.  of  6-normal  sul- 
phuric acid  and  3  cc.  of  pure  hydrofluoric  acid.  (This  acid 
should  not  be  allowed  to  come  in  contact  with  the  skin,  as  it 
produces  painful  wounds.)  Evaporate  as  far  as  possible  on  the 
steam  bath  in  a  well-drawing  hood,  adding  more  hydrofluoric 
acid  if  any  solid  residue  remains.  Cautiously  fume  off  the  sul- 
phuric acid,  heat  to  low  redness,  and  finally  ignite  over  the  blast 
lamp,  for  5-minute  periods,  to  constant  weight.  Deduct  the 
weight  of  this  residue  from  that  of  the  impure  silica,  and  from 
the  difference  calculate  the  percentage  of  SiO2  in  the  sample. 


GRAVIMETRIC  ANALYSIS^  83 

NOTES.  —  i.  The  whole  of  the  sample  must  be  ground  very  fine,  or  the 
coarser  particles  will  resist  the  action  of  the  flux ;  unless  all  of  the  material 
is  passed  through  the  bolting  cloth,  the  sifted  portion  may  not  represent 
an  average  sample. 

2.  Upon  fusion  with  sodium  carbonate,  silicates  are  decomposed  with 
the  evolution  of  carbon  dioxide.    The  other  products  of  the  decomposition 
are  sodium  silicate  and  aluminate,  ferrous  carbonate  or  ferric  oxide,  cal- 
cium and  magnesium  carbonates,  etc.    Owing  to  the  evolution  of  gas  during 
fusion,  the  heating  should  be  gradual  and  the  crucible  should  be  kept  covered. 

3.  Upon  disintegrating  the  mass  with  a  considerable  volume  of  dilute 
acid  the  silicic  acid  at  first  largely  enters  the  solution,  but  upon  evaporation 
it  is  rendered  almost  insoluble.    Treatment  of  the  fused  mass  with  strong 
acid  would  be  likely  to  cause  the  separation  of  gelatinous  silicic  acid  which 
would  inclose  metallic  salts  and  withhold  them  from  the  solution. 

4.  A  gritty  residue  remaining  after  the  disintegration  of  the  melt  with 
acid  indicates  that  the  original  silicate  has  been  but  imperfectly  decom- 
posed.   In  such  a  case  the  fusion  should  be  repeated  with  another  sample, 
which  should  be  sufficiently  well  ground  and  thoroughly  mixed  with  the  flux. 

5.  Silicic  acid  cannot  be  rendered  wholly  insoluble  by  a  single  evapora- 
tion and  heating;  nor  are  repeated  evaporations  and  moistenings  before 
filtration  as  effective  in  separating  the  silica  as  are  alternate  evaporations 
and  filtrations.    The  underlying  causes  are  as  yet  obscure. 

6.  To  free  the  silica  as  far  as  possible  from  mineral  salts,  the  residue  after 
evaporation  should  be  thoroughly  extracted  with  warm  hydrochloric  acid ; 
and  the  solution  should  be  diluted  to  a  large  volume  to  prevent  the  inclosure 
of  impurities  by  the  silica.    The  silica  is  first  washed  with  dilute  acid,  to 
prevent  the  partial  separation  of  basic  salts  of  iron,  aluminum,  etc.,  by 
hydrolysis ;  the  washing  is  then  completed  with  hot  water. 

7.  The  finely  divided  silica  holds  moisture  so  tenaciously  that  prolonged 
ignition  over  the  blast  lamp  is  necessary.    Even  then  the  ignited  powder 
tends  to  absorb  moisture,  and  it  should  therefore  be  weighed  as  rapidly  as 
possible. 

8.  Notwithstanding  all  the  precautions,  the  ignited  silica  is  rarely  pure. 
Upon  evaporation  with  hydrofluoric  and  sulphuric  acids,  however,  the 
silica  is  volatilized  as  silicon  tetrafluoride  and  water,  and  a  sulphate  residue 
is  left.    If  the  contaminating  substance  is  an  alkali  salt,  as  sodium  chloride, 
the  residue  will  remain  as  sulphate,  even  at  high  temperatures ;  but  certain 
other  sulphates,  as  those  of  iron,  aluminum,  and  titanium,  evolve  sulphur 
trioxide  on  ignition  and  leave  the  corresponding  oxides.    In  the  estimation 
of  silica,  the  weight  of  impurities  in  the  silica  is  always  determined  by 
weighing  the  residue  from  the  hydrofluoric  and  sulphuric  acid  treatment; 


84    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

in  order  then  that  the  impurities  weighed  with  the  silica  may  be  as  nearly 
as  possible  identical  with  the  residue  from  this  treatment,  it  is  best  to  treat 
the  silica  before  ignition  with  a  few  drops  of  sulphuric  acid.  The  final 
residue  from  the  hydrofluoric  and  sulphuric  acids  should  also  be  subjected 
to  the  same  temperature  employed  in  the  ignition  of  the  silica. 

9.  The  procedure  for  the  determination  in  the  united  filtrates  from  the 
silica  of  the  mixed  oxides  of  iron,  aluminum,  etc.  and  of  calcium  and  mag- 
nesium, does  not  differ  materially  from  that  given  under  the  determination 
of  calcium  and  magnesium  in  limestones. 

10.  For  a  thorough  study  of  the  analysis  of  silicate  and  carbonate  rocks, 
the  student  is  referred  to  Bulletin  No.  422  of  the  United  States  Geological 
Survey,  by  W.  F.  Hillebrand. 

THE  DETERMINATION  OF  POTASH  IN  SOLUBLE   SALTS 

The  sample  may  be  a  pure  salt,  a  soluble  industrial  product, 
or  an  artificial  mixture  of  potassium  chloride  and  sodium  car- 
bonate. 

Principle.  The  determination  of  potassium  by  this  method 
depends  upon  the  insolubility  of  potassium  per  chlorate,  and  the 
solubility  of  sodium  and  certain  other  perchlorates  in  96% 
alcohol.  This  is  not  a  precipitation  method,  but  one  of  ex- 
traction. If  heavy  metals  are  present,  they  are  first  removed. 

Procedure.  Weigh  out  samples  sufficient  to  contain  about 
0.25  g.  of  K20,  into  50  cc.  beakers,  and  treat  each  as  follows : 
Warm  the  sample  with  25  cc.  of  water,  and,  if  sulphates  are 
absent,  filter  into  a  loo-cc.  porcelain  dish ;  if,  however,  sulphates 
are  present,  acidify  with  6-normal  hydrochloric  acid,  stir,  treat 
the  hot  acid  liquid  with  barium  chloride  solution  in  slight  excess, 
filter  into  a  loo-cc.  porcelain  dish,  evaporate  the  filtrate  to  dryness 
on  the  steam  bath,  and  warm  the  residue  with  15-20  cc.  of  water, 
with  stirring.  Add  to  the  solution  sufficient  perchloric  acid  to 
contain  1.7  times  the  sample's  weight  of  HC104  (See  Note  2), 
and  evaporate  to  a  sirupy  consistency.  Add  15  cc.  of  hot  water 
and  2  cc.  of  perchloric  acid,  and  again  evaporate.  Once  more 
add  15  cc.  of  hot  water,  and  evaporate  until  heavy  fumes  of 
perchloric  acid  appear. 


GRAVIMETRIC  ANALYSIS  85 

Allow  the  mixture  to  cool  thoroughly,  add  20  cc.  alcohol  con- 
taining 0.2%  by  weight  of  HC1O4,1  and  stir  for  some  time, 
keeping  the  salt  as  coarsely  granular  as  possible.  Let  settle, 
decant  the  liquid  through  a  weighed  Gooch  crucible 2  (containing 
a  mat  moistened  with  the  wash  liquid),  and  to  the  residue  add 
a  second  20-cc.  portion  of  the  wash  liquid.  Stir,  let  settle,  again 
decant,  and  then  drive  off  the  remaining  alcohol  on  the  steam 
bath.  Dissolve  the  residue  in  15  cc.  of  hot  water,  add  a  few 
drops  of  perchloric  acid,  and  evaporate  to  heavy  fumes.  Cool, 
add  i  cc.  of  the  wash  liquid,  decant,  and  test  a  few  drops  of  the 
washings  for  complete  extraction.  (The  extraction  with  the 
alcoholic  liquid  must  be  continued  until  a  few  drops  of  the  filtrate 
leave  no  residue  when  evaporated  to  dryness  on  platinum  foil.) 
Finally,  cool,  add  i  cc.  of  the  wash  liquid,  and  sweep  the  salt 
into  the  Gooch  crucible  with  a  policeman,  washing  at  last  with 
a  very  little  pure  96%  alcohol.  Dry  the  salt  for  half  an  hour 
at  130°,  and  weigh. 

Report  the  percentage  of  K20  in  the  sample. 

NOTES.  —  i.  This  method  was  proposed  in  1831  by  Serullas,  but,  owing 
to  some  mistaken  ideas  concerning  the  properties  of  perchloric  acid,  the 
proposition  did  not  receive  the  attention  it  deserved.  Perchloric  acid 
solutions  of  satisfactory  grade  can  now  be  obtained  in  the  market,  they 
can  be  kept  indefinitely  in  glass-stoppered  bottles,  and  the  method  rivals 
in  results  the  chloroplatinic  acid  process ;  and  this  at  a  greatly  reduced  cost. 

2.  The  specific  gravities  of  perchloric  acid  solutions  are  as  follows: 
70%  HC104, 1.67  ;  60%  HC1O4,  1.54 ;  50%  HCLO4, 1.41 ;  30%  HC104,  1.20 ; 
20%  HC104,  1. 1 2.    The  strength  of  a  solution  of  the  pure  acid  may  easily  be 
determined  by  the  dilution  of  a  known  amount  and  titration  with  sodium 
hydroxide,  with  phenolphthalein  as  indicator. 

3.  In  order  to  obtain  the  potassium  as  pure  KC1O4  by  this  method,  it  is 
essential  that  no  strong  acids  be  present,  other  than  perchloric  acid,  which 
yield  salts  insoluble  in  alcohol.     Sodium  chloride  and  sulphate  are  such 
salts,  and  it  is  therefore  necessary  to  remove  chlorides  and  sulphates  before 

1  Made  by  mixing  1.7  cc.  of  the  60%  acid,  or  4.4  cc.  of  the  30%  acid,  with  one 
liter  of  96%  alcohol. 

2  It  is  better  to  use  a  Munroe  crucible,  with  a  filter  of  platinum  sponge.    The 
crucible  itself  may  be  of  gold,  to  save  expense. 


86    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  treatment  with  alcohol.  HC1  may  be  removed  by  repeatedly  evaporat- 
ing the  aqueous  solution  with  the  less  volatile  HC1O4;  but  H2S04  is  less 
volatile  than  HC1O4,  and  cannot  be  expelled  in  this  way.  Before  the  first 
evaporation,  therefore,  the  latter  should  be  precipitated  from  the  hot  acid 
solution  by  means  of  BaCl2  in  slight  excess.  Phosphates,  though  often 
insoluble  in  alcohol,  need  not  be  removed ;  but,  in  their  presence,  a  larger 
excess  of  HC1O4  should  be  used,  to  insure  their  complete  removal  by  the 
wash  liquid,  as  H3P04.  (See  Note  5  of  this  procedure,  and  also  Note  6 
under  the  determination  of  calcium.) 

4.  Since  NH4C104  is  only  sparingly  soluble  in  alcohol,  ammonium  salts 
should  be  carefully  expelled  by  gentle  ignition,  before  the  treatment  with 
HC1O4.    Moderate  amounts  of  barium,  calcium,  and  magnesium  do  not 
interfere  with  the  procedure ;  their  perchlorates  are  soluble  in  alcohol. 

5.  The  perchlorate  mixture  is  extracted  with  alcohol  containing  a  small 
amount  of  HC104  because,  owing  to  the  common  ion  effect,  the  solubility  of 
KC104  is  less  in  it  than  in  pure  alcohol ;  the  two  solubilities  are  about  4  mg. 
and  16  mg.  per  100  cc.,  respectively.    The  solubility  of  KC104  is  still  less 
in  the  presence  of  sodium,  and  other  soluble  perchlorates ;  i.e.  in  the  first 
portions  of  the  alcoholic  extract.    Alkali  and  alkali-earth  phosphates  are 
decomposed  by  perchloric  acid  and  the  H3P04  dissolves  in  the  acid-alcoholic 
liquid ;  in  the  presence  of  phosphates,  therefore,  perchloric  acid  should  be 
present  in  considerable  excess. 

6.  If  a  known  weight  of  NaCl— KC1  mixture,  obtained  for  example  in  a 
silicate  analysis,  is  converted  into  a  mixture  of  the  perchlorates,  and  the 
KC104  isolated  and  weighed,  the  method  yields  both  the  K20  and  Na2O 
contents  of  the  original  sample.     (Cf.  Part  IV,  Problem  29.) 

7.  In  order  to  prevent  the  loss  of  KC104,  in  the  separation  of  sodium 
and  potassium,  it  has  been  suggested  to  extract  the  perchlorate  mixture 
with  an  alcoholic  liquid  which  has  previously  been  saturated  with  KC104. 
This  procedure,  however,  is  apt  to  lead  to  high  results,  owing  to  the  precipi- 
tation of  small  amounts  of  potassium  from  the  wash  liquid  by  the  NaC104 
entering  into  solution ;  it  is  therefore  not  to  be  recommended.    In  order  to 
obtain  exact  results,  it  suffices  to  avoid  the  use  of  unnecessary  quantities 
of  the  wash  liquid. 

THE  ELECTROLYTIC  DETERMINATION  OF  COPPER 

The  sample  to  be  analyzed  may  be  pure  copper  sulphate,  an 
artificial  mixture  of  the  carbonates  of  copper  and  sodium,  a 
copper  ore,  or  a  nickel  coin.  In  case  stationary  electrodes  are 
employed,  the  solution  should  contain  not  over  0.2  g.  of  copper 


GRAVIMETRIC  ANALYSIS  87 

and  5  cc.  of  nitric  acid  (sp.  gr.,  1.42),  and  should  have  a  volume 
of  100  cc. ;  in  the  case  of  a  rotating  anode,  however,  the  solution 
may  contain  as  much  as  0.5  g.  of  copper,  and  it  should  contain, 
in  a  volume  of  100  cc.,  3-5  cc.  of  nitric  acid  (sp.  gr.,  1.42),  or  i  cc. 
of  sulphuric  acid  (sp.  gr.,  1.84)  and  3  g.  of  ammonium  sulphate. 

Method.  The  copper  salt  is  decomposed  by  the  electric 
current,  and  the  copper  deposited  upon  the  cathode  (negative 
electrode).  The  cathode  is  weighed  before  and  after  the  opera- 
tion, and  the  increase  in  weight  indicates  the  quantity  of  copper 
in  the  sample. 

The  polarity  of  the  terminals  may  be  determined  by  bringing 
the  wires,  about  0.5  cm.  apart,  into  contact  with  a  piece  of  filter 
paper  moistened  with  potassium  iodide  solution.  At  the  positive 
terminal  iodine  will  separate  and  color  the  paper. 

Cleaning  the  Platinum  Electrodes.1  The  electrodes  are  freed 
from  grease  by  heating  with  dilute  sodium  hydroxide  solution, 
after  which  they  are  washed  with  water;  the  cathode  is  then 
dried,  allowed  to  cool  in  a  desiccator,  and  weighed.  To  clean 
the  platinum  cathode  after  the  determination,  cover  the  deposit 
completely  with  6-normal  nitric  acid,  heat  for  at  least  15  minutes, 
and  wash. 

A.  Procedure  with  Stationary  Electrodes.2  Dissolve  a  0.5- 
o.6-g.  sample  of  copper  sulphate,  CuS04 .  5  H^O,  in  50  cc.  of 
water,  in  a  tall  150-0:.  beaker;  stir  to  complete  solution,  add 
4  cc.  of  nitric  acid  (sp.  gr.,  1.42),  and  dilute  to  100  cc.  Im- 
merse the  electrodes  in  the  solution  and  connect  them  in  such 
a  way  that  the  electrode  with  the  larger  surface  is  made  the 
cathode.  The  electrolysis,  which  should  be  carried  out  at  a 
potential  of  1.9-2.0  volts,  may  be  completed  in  the  cold  over 
night,  or  in  two  or  three  hours  if  the  temperature  is  kept  at 
70-80°  by  means  of  a  heated  sheet  of  wire  gauze  placed  a  short 
distance  below  the  beaker.  Finally  test  for  complete  deposition 
by  adding  a  little  water  to  raise  the  level  of  the  solution  on  the 

1  In  case  a  silver  cathode  is  used,  see  Note  12. 

2  If  the  sample  is  an  ore,  see  Note  n. 


88    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

cathode;  if  after  30  minutes  no  copper  is  to  be  seen  upon  the 
fresh  platinum  surface,  the  deposition  is  probably  complete. 
(Test  a  few  cubic  centimeters  of  the  solution  with  sodium  acetate 
and  a  drop  of  potassium  ferrocyanide  solution.)  Without  dis- 
connecting the  electrodes,  siphon  off  the  electrolyte  while  in- 
troducing distilled  water,  until  the  current  ceases  to  pass ;  this 
is  to  prevent  the  re-solution  of  any  of  the  copper  by  the  acid 
liquid.  Remove  the  cathode,  wash  it  with  water,  then  with 
alcohol,  and  dry  it  for  a  short  time  in  an  air  bath  at  85-90°. 
Allow  the  cathode  to  cool  in  a  desiccator,  and  weigh. 

B.  Procedure  with  a  Rotating  Anode.1  Heat  a  five-cent  coin 
with  sodium  hydroxide  solution  to  free  it  from  grease,  then  wash 
it  with  water,  and  dry  at  100°.  After  cooling,  weigh  the  coin, 
and  dissolve  it  in  50  cc.  of  6-normal  nitric  acid,  in  a  covered 
casserole.  Evaporate  the  solution  to  dryness  on  the  steam 
bath,  dissolve  the  residue  in  about  100  cc.  of  cold  water,  add 
10  cc.  of  sulphuric  acid  (sp.  gr.,  1.84),  allow  to  cool,  and  transfer 
the  whole  to  a  5oo-cc.  measuring  flask,  diluting  to  the  mark  with 
water.  Measure  off  one  tenth  of  the  well-mixed  solution  into 
a  5o-cc.  graduated  flask  and  transfer  this  quantitatively  to  the 
electrolytic  vessel;  before  transferring  the  solution  to  the  elec- 
trolytic vessel,  however,  pour  100  cc.  of  water  into  the  latter 
and  adjust  the  electrodes  so  that,  when  four  fifths  covered  with 
water,  they  do  not  come  into  contact  with  one  another,  nor 
cause  a  loss  of  liquid,  when  the  anode  is  rotated.  The  water 
can  then  be  siphoned,  or  drawn  off,  and  the  solution  transferred 
to  the  vessel  and  diluted  to  100  cc.  without  disturbing  the  vessel 
or  the  electrodes. 

To  perform  the  electrolysis,  attach  the  anode  to  the  shaft 
of  the  rotator  (which  is  connected  by  means  of  a  mercury  cup, 
or  otherwise,  with  the  positive  terminal),  and  the  cathode  to 
the  negative  terminal.  Adjust  the  levels  so  that,  with  100  cc. 
of  solution,  the  cylindrical  cathode  is  about  four  fifths  im- 
mersed in  the  liquid.  Then,  by  means  of  the  sliding  contact, 
1  If  the  sample  is  an  ore,  see  Note  u. 


GRAVIMETRIC  ANALYSIS  89 

throw  in  the  maximum  resistance  of  the  rheostat,  see  that  all 
connections  are  well  made,  start  the  motor,  and  close  the  switch ; 
immediately  decrease  the  resistance  of  the  rheostat  until  the 
ammeter  registers  about  i  ampere,  and  allow  the  electrolysis 
to  proceed.  In  the  presence  of  nickel,  the  voltage  should  not  ex- 
ceed 2.7.  After  about  55  minutes,  test  for  complete  deposition 
by  adding  a  little  water  to  raise  the  level  of  the  solution  on  the 
cathode ;  if  after  10  minutes  no  copper  is  visible  on  the  freshly 
exposed  platinum,  the  deposition  is  complete.  When  this  is 
the  case,  without  disconnecting  the  terminals,  stop  the  rotator 
and  draw  off  the  solution  into  a  large  beaker,  carefully  pouring 
in  water  as  fast  as  the  solution  flows  out.1  As  soon  as  the 
ammeter  indicates  that  no  current  is  passing,  throw  off  the  switch, 
remove  the  cathode  and  wash  off  the  water  with  a  little  alcohol ; 
dry  below  100°,  allow  to  cool  in  a  desiccator,  and  weigh. 

NOTES.  —  i.  If  two  platinum  plates,  immersed  in  an  aqueous  solution 
of  copper  sulphate,  are  connected  by  wire  with  the  poles  of  a  storage  battery, 
metallic  copper  will  be  deposited  upon  one  of  the  plates;  under  certain 
conditions,  all  of  the  copper  will  separate  in  the  form  of  a  compact,  firmly 
adherent  metallic  film. 

The  process  of  decomposition  is  called  electrolysis;  the  solution  under- 
going decomposition  is  called  an  electrolyte;  the  two  poles  by  which  the 
current  enters  and  leaves  the  electrolyte  are  called  electrodes.  When  salt 
solutions  are  electrolyzed,  the  positive  ions  (cations)  move  towards  the 
negative  electrode  (cathode),  and  the  negative  ions  (anions)  towards  the 
positive  electrode  (anode). 

The  quantity  of  electricity  which  passes  through  the  solution  in  unit 
tune,  or  the  speed  of  the  current,  is  measured  by  an  ammeter.  The  unit, 

1  If  it  is  desired  to  determine  the  nickel  electrolytically,  evaporate  this  dilute 
solution  to  a  volume  of  25-30  cc.,  make  slightly  alkaline  with  ammonia,  filtering 
off  any  ferric  hydroxide  which  may  be  precipitated,  and  to  the  solution  (40  cc.  in 
volume)  in  the  electrolytic  vessel  add  60  cc.  of  ammonia  of  sp.  gr.  0.90.  Elec- 
trolyze  at  3.0-3.5  volts  with  a  rotating  anode.  After  about  an  hour,  test  for  com- 
plete deposition  by  adding  to  a  few  drops  of  the  solution,  neutralized  with  acetic 
acid,  a  drop  or  two  of  dimethyl-glyoxime  solution  (a  red  color  indicates  nickel). 
When  the  deposition  is  complete,  proceed  as  directed  in  the  copper  determination. 
Finally  remove  the  nickel  from  the  cathode  by  heating  for  at  least  15  minutes  with 
6-normal  nitric  acid. 


QO    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

called  an  ampere,  is  represented  by  the  unvarying  current  which,  when 
passed  through  a  solution  of  silver  nitrate,  deposits  metallic  silver  at  the 
rate  of  0.001118  g.  per  second.1 

The  electromotive  force,  i.e.  the  electrical  pressure  which  drives  the 
current  along  the  circuit,  is  measured  by  the  voltmeter.  The  unit,  called 
a  wit,  is  represented  by  the  electrical  pressure  that  produces  a  current  of 
one  ampere  when  steadily  applied  to  a  conductor  whose  resistance  is  one 
ohm. 

The  unit  of  resistance,  called  an  ohm,  is  represented  by  the  resistance 
offered  to  an  unvarying  electric  current  by  a  column  qf  mercury  14.4521  g. 
in  mass,  of  a  constant  sectional  area  and  a  length  of  106.3  cm->  at  the  tem- 
perature of  melting  ice. 

These  magnitudes  are  always  related  to  one  another  as  follows  (Ohm's 
law): 

Quantity  of  electricity  (amperes)  =  Electromotive  force  (volts)  Qr  .=£ 

Resistance  (ohms)  R 

The  most  satisfactory  current  producer  for  electro-analysis,  in  which  a 
steady  non-fluctuating  current  is  desired,  is  the  secondary  or  storage  ele- 
ment. The  E.  M.  F.  of  the  lead  cell  is  about  2  volts,  and  the  necessary 
voltage  for  the  work  may  be  obtained  by  connecting  several  cells  in  series. 
In  practice,  the  potential  difference  between  the  electrodes  is  regulated  by 
means  of  incandescent  lamps,  coils  of  wire,  or  other  devices,  which  offer 
resistance  to  the  flow  of  electricity  along  the  circuit  and  convert  electrical 
energy  into  heat.  Any  rheostat  will  do  for  this  work,  provided  the  range 
in  resistance  is  properly  related  to  the  other  factors  which  determine  the 
current  strength.  The  sliding-contact  coil  resistances  which  are  on  the 
market  are  very  satisfactory.  Voltmeters  and  ammeters  should  have  the 
scales  graduated  with  a  range  as  limited  as  is  consistent  with  the  current 
conditions  to  be  employed,  so  that  each  subdivision  may  represent  a  small 
fraction  of  a  unit.  The  manner  of  connecting  the  instruments  is  illus- 
trated in  the  accompanying  figure. 

2.  The  passage  of  an  electric  current  of  suitable  voltage  through  the 
solution  of  an  ionogen  is  associated  with  physical  and  chemical  changes 
which  often  may  be  utilized  in  exact  gravimetric  analysis. 

1  The  quantity  of  a  given  metal  deposited  by  a  current  of  electricity  is  directly 
proportional  to  the  quantity  of  electricity  which  passes  through  the  solution; 
and  the  quantities  of  different  metals  deposited  by  a  specific  quantity  of  electricity 
are  directly  proportional  to  the  chemical  equivalents  of  the  metals  in  the  solution. 
These  two  statements  are  known  as  Faraday's  laws,  though  these  apply  to  non- 
metallic  ions  as  well. 


GRAVIMETRIC  ANALYSIS 


91 


The  chemical  effect  at  the  cathode  is  always  some  form  of  reduction. 
Simple  metallic  ions,  as  those  of  copper,  tin,  nickel,  cobalt,  cadmium,  etc., 
travel  towards  the  cathode,  where  they  give  up  their  charges  and  separate 


Electrodes 


in  the  metallic  condition;  while  the  hydrogen  ion  here  loses  its  positive 
charge  and  either  acts  directly  as  a  reducing  agent  (e.g.  nitric  acid  to  am- 
monia) or  is  evolved  as  gaseous  hydrogen. 

At  the  anode,  on  the  other  hand,  the  chemical  effect  is  always  some  form 
of  oxidation.  The  anions  of  the  halogen  group  are  liberated  as  free  chlorine, 
bromine,  or  iodine  and  may  act  as  oxidizing  agents,  while  from  solutions 
containing  hydroxide,  sulphate,  or  nitrate  ions,  oxygen  separates  at  the 
anode  and  either  acts  directly  as  an  oxidizing  agent  or  is  evolved  in  gaseous 
form.  (It  should  be  borne  in  mind  in  this  connection  that  aqueous  solutions 
always  contain  the  ions  of  water.)  Although  positive  ions  always  move 
towards  the  cathode,  certain  metals  (e.g.  lead,  cobalt,  nickel,  and  a  few 
others)  may,  under  specific  conditions,  be  oxidized  (possibly  to  complex 
oxy-anions)  and  deposited  more  or  less  completely  at  the  anode  in  the  form 
of  insoluble  peroxides.  In  fact,  lead  can  be  determined  accurately  in  this 
way,  as  an  oxide. 

3.  Metals,  like  all  other  substances,  possess  when  immersed  in  water  a 
characteristic  solution  tension,  by  which  is  understood  an  expansive  force 
which  seeks  to  drive  particles  of  the  metal  into  the  solution  ;  when  a  metal 
is  immersed  in  the  solution  of  one  of  its  salts  it  will  either  send  more  of  its 
atoms  into  the  solution  as  ions,  or  some  of  its  ions  will  be  discharged  from 
the  solution  on  its  surface  as  atoms.  In  the  first  case  the  metal  will  be- 
come negatively  charged,  and  in  the  second  case  positively  charged  with 
respect  to  the  solution  ;  in  either  case  equilibrium  will  be  reached  when 
the  solution  tension  of  the  metal  is  exactly  counterbalanced  by  the 
electrostatic  charges  and  the  osmotic  pressure  of  the  metallic  ions  in  the 
solution. 

Upon  comparing  the  different  elements  from  the  standpoint  of  the 
potential  difference  between  them  and  their  salt  solutions,  at  identical 
normal  ion-concentrations,  a  characteristic  series  of  values  is  obtained. 


92     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

In  the  case  of  solutions  of  normal  ion  concentration,  for  example,  some  of 
the  values  are  as  follows : 

Zn= +0.493  Ag=  — 1.05 

Cd= +0.143  I  = -0.80 

Fe= +0.067  Br=  — 1.27 

Co=  —0.045  O=  — 1.50    (At  N.  H+-ion  concentration.) 

Ni=  —0.049  Cl=  — 1.63 

H= -0.277  S04=-2.i8 
Cu=- 0.606 

4.  In  the  electrolysis  of  a  salt  solution,  both  electrodes  soon  become 
coated  with  the  products  deposited  (i.e.  each  becomes  essentially  an  elec- 
trode of  the  deposited  material,  no  matter  what  its  original  composition) 
and  they  are  said  to  be  polarized.  Hence  a  system  similar  to  that  just  dis- 
cussed may  be  considered  to  exist  at  each  electrode,  and  it  is  evident  that 
electrolysis  must  act  in  opposition  to  the  solution  tension  of  the  elements 
and  in  conjunction  with  the  osmotic  pressure  of  their  ions.  In  order  then 
to  decompose  a  salt  solution  continuously,  a  voltage  at  least  slightly  in 
excess  of  the  polarization  voltage  must  be  applied;  i.e.  a  voltage  greater 
than  the  numerical  difference  between  the  single  potential  differences 
normally  established  at  the  cathode  and  anode.  Assuming  normal  ion- 
concentrations,  the  decomposition  voltage  of  copper  chloride,  for  example, 
is  —1.63  minus  —0.606=1.02  volts;  but  the  decomposition  voltage  of  a 
solution  as  calculated  in  this  manner,  especially  in  the  case  of  a  salt  of  an 
oxyacid,  frequently  fails  to  agree  with  that  found  by  experiment.  The 
separation  of  gases  at  the  electrodes  is  often  accompanied  by  "overvoltages," 
which  vary  more  or  less  markedly  with  the  material  and  physical  nature  of 
the  electrodes ;  moreover,  the  ion-concentration  is  generally  unknown,  and 
it  always  changes  as  the  electrolysis  proceeds.  The  important  matter 
here  lies  not  so  much  in  the  calculation  of  decomposition  voltages  as  in  the 
recognition  of  the  existence  of  a  minimum  decomposition  voltage  for  every 
ionogen,  under  definite  conditions. 

In  the  case  of  an  ionogen  of  known  decomposition  voltage,  we  should 
simply  use  a  somewhat  higher  voltage,  but  if  other  metallic  ions  were 
present  it  might  be  impossible  to  completely  deposit  one  metal  without 
using  a  voltage  that  would  start  the  deposition  of  the  second  metal  also. 
While  copper  can  readily  be  separated  from  cobalt  or  from  nickel  by  elec- 
trolysis, it  is  not  possible  to  separate  nickel  from  cobalt  in  this  way ;  and  in 
general  only  metals  whose  deposition  voltages  differ  by  several  tenths  of  a 
unit  can  be  separated  from  each  other  by  maintaining  an  intermediate 
voltage  during  the  electrolysis. 


GRAVIMETRIC  ANALYSIS  93 

The  addition  of  certain  reagents,  as  ammonia,  potassium  cyanide,  am- 
monium oxalate,  etc.,  to  solutions  containing  two  metals  sometimes  reduces 
the  concentration  of  one  metallic  ion  very  much  more  than  that  of  the 
other,  owing  to  the  formation  of  more  or  less  stable  complexes,  and  makes 
it  possible  to  perform  a  separation  by  the  "constant  voltage"  method  that 
otherwise  might  not  be  possible. 

5.  The  current  strength  will  of  course  depend  upon  the  voltage  used, 

since,  according  to  Ohm's  law,  i  =  — .    In  performing  an  electrolysis,  the 

R 

voltage  actually  available  is  diminished  by  the  decomposition  voltage  of 
the  electrolyte  (polarization  voltage) ;  hence  the  current  which  passes  is 
equal  to  the  available  voltage  minus  the  decomposition  voltage  of  the 
electrolyte,  divided  by  the  resistance  of  the  circuit. 

6.  The  quantity  of  metal  deposited  in  a  given  time  is  dependent  upon 
the  strength  of  the  current  in  amperes.    A  current  of  i  ampere  is  capable 
of  depositing  1.118  mg.  of  silver,  and,  according  to  Faraday's  law,  equiva- 
lent amounts  of  other  elements,  per  second.    This  law  might  be  used  to 
calculate  the  time  necessary  for  the  complete  deposition  of  the  metal  if, 
under  the  analytical  conditions,  it  were  the  only  cation  taking  part  hi  the 
electrolysis.    In  the  neighborhood  of  the  cathode,  however,  the  concentra- 
tion of  the  solution  with  respect  to  this  cation  gradually  decreases  to  an 
infinitesimal  value,  and  the  resistance  and  the  decomposition  voltage  of  the 
solution  therefore  rise ;  finally  a  point  is  reached  at  which  other  ions  begin 
to  be  discharged.     Since  circulation  of  the  solution  tends  to  maintain  a 
uniform  distribution  of  the  ions,  mechanical  stirring  favors  the  rapid  deposi- 
tion of  those  ions  which  have  the  lowest  discharge  voltages. 

7.  Unless  the  solution  is  mechanically  stirred,  the  rate  of  deposition  of 
a  given  metal  decreases  rapidly,  owing  to  the  decreasing  concentration  of 
its  ions  around  the  cathode  and  to  the  continually  increasing  proportion 
of  the  current  which  is  carried  by  the  hydrogen  (or  other)  ions.     Since  a 
rapid  circulation  of  the  solution  tends  greatly  to  prevent  the  local  decrease 
in  metallic  ion  concentration  around  the  cathode,  and  since  with  improved 
circulation  currents  of  much  higher  density  may  be  used  than  would  other- 
wise give  satisfactory  deposits,  it  is  possible  to  greatly  reduce  the  time  neces- 
sary for  a  determination  by  performing  the  electrolysis  with  the  use  of  a 
rotating  electrode. 

8.  Owing  to  the  reduced  viscosity  at  higher  temperatures,  the  resistance 
offered  by  an  aqueous  solution  to  the  passage  of  electricity  decreases  with 
a  rise  in  temperature,  and  in  this  way  the  voltage  required  to  produce  a 
given  current  may  be  reduced  to  a  minimum;   this  may  sometimes  be 
of  importance  in  electrolytic  separations.     Moreover,  in  case  stationary 


94    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

electrodes  are  used,  heating  the  solution  during  electrolysis  gives  rise  to 
more  or  less  rapid  convection  currents,  and  also  increases  the  speed  of 
diffusion,  and  these  effects  are  equivalent  to  a  gentle  mechanical  stirring. 
The  solution  should  never  be  heated  to  the  boiling  point,  however,  since  the 
deposit  might  in  that  case  be  loosened  from  the  cathode. 

9.  The  deposited  metal  tends  to  redissolve  in  the  electrolyte  (cf.  "polar- 
ization"), and  consequently  the  rate  at  which  the  metal  is  deposited  must 
exceed  that  at  which  it  redissolves.    The  metal  is  only  deposited  from  the 
solution  in  immediate  contact  with  the  cathode,  so  that  the  greater  the  area 
of  the  cathode,  the  more  metal  there  is  available  for  deposition ;  but  also 
the  greater  the  rate  of  re-solution.    Hence  it  follows  that  the  current 
strength  necessary  for  the  satisfactory  deposition  of  the  metal  is  propor- 
tional to  the  area  of  the  cathode. 

The  current  strength  per  unit  area  is  called  the  current  density;  a  square 
decimeter  is  generally  taken  as  the  unit  area.  Hence  a  "normal  current 
density  of  2  amperes  "  means  a  current  of  2  amperes  per  100  sq.  cm.  of  cathode 
area,  or  of  i  ampere  for  50  sq.  cm.  of  cathode  area,  etc.  While  the  tendency 
of  a  metal  to  redissolve  fixes  a  lower  limit  for  the  current  density  to  be  used, 
a  higher  limit  is  set  by  the  tendency  of  the  metal  to  form  spongy,  non- 
adherent  films  when  deposited  too  rapidly. 

It  is  highly  important  for  accurate  work  to  deposit  the  metal  in  the 
form  of  a  compact  film  which  can  easily  be  washed  and  weighed  without 
loss.  The  condition  of  the  deposit  depends  not  only  upon  the  current 
density  used,  but  also  upon  the  concentration  of  the  metallic  ions  in  the 
solution,  the  amount  of  free  acid  and  other  substances  present,  the  tem- 
perature, etc.  The  best  conditions  for  specific  cases  have  been  determined 
by  repeated  experiments. 

10.  Concerning  the  effect  upon  the  nature  of  the  deposit  of  the  products 
that  accumulate  in  the  solution  during  an  electrolysis,  that  are  purposely 
added  to  the  solution,  or  which  were  originally  present  in  the  sample,  it 
may  be  stated  that  a  very  marked  influence  is  often  exerted  by  certain 
acids,  bases,  and  other  substances.    A  solution  of  copper  sulphate,  if 
electrolyzed  without  the  addition  of  another  substance,  is  almost  sure  to 
give  a  reddish  brown,  non-adherent  deposit  of  spongy  copper ;  the  addi- 
tion of  a  little  sulphuric  acid  gives  rise  to  a  much  more  compact  deposit, 
while  the  addition  of  nitric  acid  leads  to  a  still  better  deposit  of  bright  red 
firmly  adherent  metal.    A  small  quantity  of  urea,  in  addition  to  either  acid, 
appears  to  favor  still  more  the  formation  of  a  satisfactory  deposit.    On  the 
other  hand,  high  current  densities,  which  cause  a  rapid  discharge  of  hydro- 
gen, are  apt  to  yield  loosely  adherent  deposits  of  spongy  metal.    A  current 
density  which  gives  a  bright  red,  coherent  deposit  of  pure  copper  when  no 


GRAVIMETRIC  ANALYSIS  95 

interfering  substance  is  present,  will  often  give  a  very  dark,  loosely  adherent 
deposit  when  arsenic  is  present,  even  in  small  amount.  Such  impurities 
must  be  removed  before  the  electrolysis  is  begun. 

11.  If  the  sample  to  be  analyzed  by  this  method  is  a  copper  ore,  and  is 
not  known  to  be  free  from  arsenic  and  other  interfering  substances,  it  should 
be  subjected  to  special  treatment,  in  order  to  obtain  a  solution  suitable  for 
electrolysis.     In  most  cases,  a  satisfactory  solution  may  be  prepared  accord- 
ing to  the  procedure  detailed  under  the  volumetric  estimation  of  copper 
(which  see) ;  the  ore  is  evaporated  with  aqua  regia,  the  residue  extracted 
with  dilute  hydrochloric  acid  and  water,  and  the  copper,  arsenic,  etc.,  pre- 
cipitated with  sodium  thiosulphate ;  the  arsenic  is  then  driven  off  by  igni- 
tion, the  residue  evaporated  to  dryness  with  nitric  acid,  and  finally  taken 
up  in  4  cc.  of  nitric  acid  (sp.  gr.,  1.42)  and  50  cc.  of  water.    This  solution 
is  diluted  to  100  cc.  and  electrolyzed. 

12.  The  electrode  material  should  preferably  be  insoluble  in  the  electro- 
lyte, with  or  without  current  action,  and  for  that  reason  platinum  is  most 
often  used  for  electrodes ;  but  the  continued  advance  in  the  price  of  this  metal 
has  led  to  a  search  for  less  expensive  materials.    Other  metals,  as  silver  and 
copper,  are  in  some  cases  suitable  for  use  as  cathodes  in  the  deposition  of 
metals  (as  is  also  the  more  expensive  palladium-gold  alloy  which  is  in  the  mar- 
ket), but  the  anode  must  still  be  made  of  platinum  or  of  something  equally 
resistant.     In  the  determination  of  copper,  for  example,  a  silver  cathode  is 
about  as  satisfactory  as  one  of  platinum ;  the  deposit  can  be  removed  by 
means  of  dilute  hydrochloric  acid,  with  the  addition  of  a  little  hydrogen 
peroxide  or  nitric  acid,  and,  after  washing  with  ammonia,  the  cathode  is 
again  ready  for  use. 

Since  a  practical  limit  is  placed  upon  the  current  density,  the  time  neces- 
sary for  a  deposition  is  inversely  proportional  to  the  area  of  the  cathode ; 
for  this  reason  the  electrode  to  receive  the  deposit  should  present  the  maxi- 
mum of  surface  to  the  solution. 

The  platinum  dish  electrode  designed  by  Classen  is  quite  thin  and  presents 
a  relatively  large  surface ;  a  Classen  dish  weighing  40  g.  has  a  capacity  of 
about  250  cc.  and  presents  an  inner  surface  of  about  150  sq.  cm.  to  the  solu- 
tion. A  platinum  disk  or  a  flat  spiral  of  platinum  wire  may  be  used  as  the 
anode.  The  electrodes  commonly  used,  however,  are  more  economical. 
They  are  open  cylinders  of  thin  foil  or  of  fine  mesh  gauze,  and  elongated 
spirals  of  heavy  platinum  wire ;  the  cylinders,  which  are  used  to  receive  the 
deposit,  weigh  10-12  g.  and  the  wire  spirals  about  8  g.  Of  all  cathodes, 
those  of  gauze  are  the  most  efficient ;  they  present  a  relatively  larger  sur- 
face, all  parts  of  the  surface  are  equally  effective,  they  permit  a  much  better 
circulation  of  the  solution  and  consequently  the  use  of  higher  current  den- 


g6    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

sities.    These  cylindrical  cathodes  and  spiral  anodes  are  the  ones  assumed 
in  the  foregoing  procedures. 

13.  In  case  the  electrolysis  is  to  be  performed  with  stationary  electrodes, 
it  is  best  to  use  a  tall  beaker  of  small  diameter,  which  can  be  heated.  If 
heating  is  not  desired,  however,  or  if  a  rotating  electrode  is  used,  the  most 
suitable  vessel  is  a  150-0:.  glass  cylinder,  with  a  rounded  bottom  ending  in 
an  outlet  tube  provided  with  a  stopcock ;  the  electrodes  should  reach  nearly 
to  the  bottom  of  this  cylinder,  to  insure  efficient  mixing.  After  the  deposition 
is  complete,  without  interrupting  the  current,  the  electrolyte  can  easily  be 
drawn  off  with  the  simultaneous  introduction  of  distilled  water  above. 


PART   III 

VOLUMETRIC  ANALYSIS 

GENERAL  DISCUSSION 

Fundamental  Principles.  It  has  already  been  pointed  out 
in  Part  I  that  in  volumetric  analysis  the  amount  of  an  element 
or  compound  present  in  a  sample  is  calculated  from  the  volume 
of  some  reagent  of  known  concentration  which  is  required,  after 
suitable  treatment  of  the  sample,  to  complete  a  definite  reaction. 

The  analytical  balance  is  equally  requisite  as  a  starting  point 
for  both  gravimetric  and  volumetric  systems;  in  addition  to 
the  balance,  volumetric  processes  demand  graduated  measur- 
ing instruments  and  standard  solutions  (i.e.  solutions  of  ac- 
curately known  value).  The  concentration  or  value  of  a  solu- 
tion for  a  specific  reaction  is  determined  by  a  procedure  called 
standardization,  in  which  the  solution  is  brought  into  reaction 
with  a  definite  weight  of  a  substance  of  known  purity ;  from  the 
volume  of  solution  required  to  complete  the  reaction,  the  strength 
or  value  of  the  solution  can  be  calculated,  and  it  is  then  a  stand- 
ard solution. 

The  value  of  standard  solutions  may  be  expressed  in  terms 
of  the  weight  of  reagent  actually  present  in  each  cubic  centimeter, 
or,  better,  in  terms  of  the  weight  of  a  given  substance  with  which 
one  cubic  centimeter  of  the  solution  will  react;  but  since  the 
weight  of  reagent  present  in  a  unit- volume  is  always  chemically 
equivalent  to  the  weight  of  substance  with  which  the  unit- volume 
reacts,  it  is  in  general  more  convenient  to  express  the  value  of 
the  standard  solution  in  terms  of  chemical  equivalents  per  unit- 
volume.  Such  solutions  are  often  made  to  bear  some  simple 
H  97 


98    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

relation  to  a  normal  solution  of  the  specific  reagent;  they  are, 
for  example,  half-normal,  tenth-normal,  or  fiftieth-normal  solu- 
tions. A  normal  solution  contains  in  one  liter  one  gram-equiva- 
lent of  the  active  reagent;  i.e.  that  quantity  of  the  active  re- 
agent which  contains,  unites  with,  replaces,  or  in  any  way,  di- 
rectly or  indirectly,  brings  into  reaction  1.008  g.  of  hydrogen. 

Thus  a  liter  of  normal  acid  solution  will  contain  1.008  g.  of 
available  hydrogen  ion  (e.g.  one  mol  of  HC1,  or  one  half  mol  of 
H2S04,  etc.) ;  and  a  liter  of  normal  alkali  solution  will  contain 
sufficient  available  hydroxide  ion  to  combine  with  1.008  g.  of 
hydrogen  ion,  or  17.008  g.  (e.g.  one  mol  of  NaOH,  one  half  mol 
of  Ba(OH)2,  etc.).  A  normal  solution  of  an  oxidizing  agent 
will  have  the  same  oxidizing  value  per  liter  as  one  gram-equiva- 
lent, 8.00  g.,  of  oxygen  (e.g.  one  sixth  mol  of  K^C^Oj,  etc.) ; 
and  a  liter  of  normal  reducing  agent  will  have  the  same  reducing 
value  as  i  .008  g.  of  hydrogen  (e.  g.  one  half  mol  of  SnCy .  It  will 
be  seen  that  a  liter  of  normal  acid  solution  will  exactly  neutralize 
a  liter  of  normal  alkali  solution,  and  a  liter  of  normal  oxidizing 
solution  will  exactly  oxidize  a  liter  of  normal  reducing  solution, 
and  so  on. 

It  should  be  especially  noted,  however,  that  the  equivalent 
or  normal  weight  of  a  substance  may  vary  according  to  the  reaction 
in  which  it  is  used.  Thus  the  normal  weight  of  oxalic  acid  is 
one  half  its  molecular  weight,  whether  it  be  used  as  a  neutraliz- 
ing, a  reducing,  or  a  precipitating  agent;  whereas  the  normal 
weight  of  nitrous  acid  would  be  the  molecular  weight,  if  used 
either  as  a  neutralizing  agent  or  to  oxidize  hydriodic  acid,  but 
only  one  half  the  molecular  weight  if  used  to  reduce  potassium 
permanganate.  In  the  case  of  potassium  permanganate,  two 
molecules  yield  three  atoms  of  available  oxygen  (equivalent  to 
six  hydrogen  atoms)  in  neutral  solution,  and  five  atoms  of  avail- 
able oxygen  (equivalent  to  ten  atoms  of  hydrogen)  when  used 
in  acid  solution.  The  normal  weight  of  this  compound  as  an 
oxidizing  agent  is  therefore  one  third  or  one  fifth  of  the  molec- 
ular weight,  according  to  the  conditions  of  its  use. 


VOLUMETRIC  ANALYSIS  99 

The  preparation  of  exactly  normal,  half-normal,  or  tenth- 
normal  solutions  generally  requires  considerable  time  and  care, 
and  is  usually  carried  out  only  when  a  large  number  of  analyses 
are  to  be  made,  or  when  the  analyst  has  some  other  specific  pur- 
pose in  view.  It  is  much  easier  to  prepare  standard  solutions 
which  differ  but  slightly  from  half-normal  or  tenth-normal, 
and  these  still  have  the  advantage  of  approximate  equality; 
two  approximately  half-normal  solutions  are  much  more  con- 
venient to  work  with  than  two  which  are  widely  different  in 
strength.  When  these  approximate  solutions  are  used,  the 
volumes  can  readily  be  reduced  to  the  corresponding  values  in 
terms  of  solutions  which  are  exactly  normal,  half-normal,  or 
tenth-normal.  For  example,  25.75  cc-  °f  a  0.0987  N  solution 
are  equivalent  to  25.75X0.0987  =  2.542  cc.  of  the  normal,  or  to 
2.542  X 10  =  25.42  cc.  of  the  tenth-normal  solution. 

Reactions  Suitable  for  Volumetric  Processes.  Volumetric 
processes  are  usually  based  upon  definite  chemical  reactions, 
and  in  general  only  such  reactions  are  suitable  as  can  be  made 
to  take  place  completely  and  very  rapidly  when  equivalent 
amounts  of  the  reacting  substances  are  brought  together.  Vol- 
umetric determinations,  however,  do  not  always  consist  in  the 
direct  titration  of  the  substances  under  investigation.  In  many 
cases  an  excess  of  the  standard  solution  is  used,  and  this  excess 
is  then  titrated  with  a  second  standard  solution.  The  volumetric 
relation  between  the  two  standard  solutions  being  known,  the 
proper  correction  for  the  excess  of  the  first  solution  is  easily 
made.  This  is  known  as  the  method  of  back  titration.  In  other 
cases  the  substance  under  investigation  is  capable  under  suitable 
conditions  of  setting  free  or  of  carrying  down  as  a  precipitate 
a  definite  proportion  of  some  other  substance  which  can  subse- 
quently be  titrated  with  a  standard  solution ;  from  the  volume 
of  the  latter  required  it  is  an  easy  matter  to  calculate  directly 
the  weight  of  the  original  substance. 

The  processes  of  volumetric  analysis  are  readily  classified, 
according  to  their  character,  into : 


100    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

A.  Neutralization  Methods;  such  as  those  of  acidimetry  and 
alkalimetry. 

B.  Methods  of  Oxidation  and  Reduction;   as  exemplified  in 
the  determination  of  ferrous  iron  by  oxidation  with  potassium 
permanganate. 

C.  Precipitation  Methods;   as,  for  example,  the  titration  of 
silver  with  potassium  thiocyanate  solution. 

Determination  of  the  End-point.  In  order  to  utilize  a  reaction 
for  volumetric  purposes,  it  is  necessary  to  have  some  means  of 
ascertaining  the  point  at  which  an  equivalent  volume  of  the 
standard  solution  has  been  added.  In  the  neighborhood  of 
this  point  certain  physical  and  chemical  properties  of  the  solu- 
tion change  very  rapidly;  in  many  cases  there  is  a  marked 
change  in  color,  electrical  conductivity,  or  oxidation  potential 
of  the  solution,  and  in  some  a  precipitate  just  ceases  to  form. 
In  numerous  cases  the  presence  of  another  reagent,  called  an 
indicator,  gives  rise  to  a  decided  color  change  in  the  solution, 
or  in  a  few  instances  it  causes  a  precipitate  to  form  and  thus 
renders  the  solution  turbid.  The  point  at  which  the  standard 
solution  has  been  added  in  sufficient  quantity  to  make  these 
changes  apparent  is  called  the  end-point  of  the  titration.  If 
the  process  is  to  be  sufficiently  accurate,  the  difference  between 
this  point  and  the  true  end-point  of  the  reaction  (i.e.  the  point 
at  which  an  equivalent  quantity  of  the  solution  has  been  added) 
must  be  exceedingly  small;  and  this  is  actually  the  case  in 
nearly  all  of  the  established  volumetric  processes. 

In  a  few  cases,  however,  a  principle  which  can  be  used  to  great 
advantage  is  that  of  compensating  errors ;  here  the  errors  which 
are  involved  in  the  actual  determination  are  counteracted  by 
equal  errors  in  the  standardization  of  the  solution  used.  Assum- 
ing that  there  is  a  noticeable  discrepancy  between  the  true  and 
the  observed  end-point,  it  will  generally  hold  good  that  this 
discrepancy  will  remain  constant  so  long  as  the  conditions  are 
the  same.  If  the  solution  can  be  standardized  under  conditions 
identical  with  those  which  obtain  in  the  actual  determination, 


VOLUMETRIC  ANALYSIS''  ioi' 

all  errors  can  be  practically  eliminated;  the  standard  solution 
becomes  merely  an  instrument  for  comparing  two  solutions  of 
the  same  substance,  one  representing  a  known  amount  of  the 
standard  solution  and  the  other  an  approximately  equal  but 
unknown  amount.  Considered  from  this  point  of  view,  when- 
ever it  is  possible  to  use  the  same  reagent  for  the  determination 
of  different  substances,  strict  accuracy  would  demand  that  the 
solution  be  standardized  by  comparison  with  a  known  weight 
of  the  substance  for  which  it  is  to  be  used  in  a  given  case. 

General  Theory  of  Indicators.  Whenever  a  substance  is 
titrated  in  the  presence  of  an  indicator,  the  physical  change 
which  enables  us  to  recognize  the  end-point  is  the  visible  result 
of  a  chemical  reaction  in  which  the  indicator  itself  is  an  active 
reagent.  In  case  the  indicator  is  acted  upon  by  the  standard 
solution  employed  in  making  the  titration,  the  reactions  involved 
may  be  represented  by  the  following  equations,  in  which  X 
represents  the  substance  being  titrated,  R  the  reagent  in  the 
standard  solution,  and  I  the  indicator  employed : 

(1)  X+R^RX, 

(2)  I+R3ZIR, 

(3)  IR+X^RX+L 

The  appearance  of  the  end-point  is  here  dependent  upon  the 
concentration  of  IR,1  which  should  remain  equal  to  zero  as  long 
as  an  appreciable  amount  of  X  is  present,  but  should  increase 
in  direct  proportion  to  the  amount  of  R  that  is  added  after  the 
concentration  of  X  has  been  reduced  to  an  infinitesimal  value. 
In  other  words,  reaction  (i)  must  be  completed  before  reaction 
(2)  begins  to  take  place,  but  reaction  (2)  must  take  place  promptly 
from  left  to  right,  even  at  an  exceedingly  low  concentration  of 

1  IR  is  not  necessarily  a  compound  of  the  indicator  with  R,  but  even  so  a  definite 
concentration  of  the  free  reagent  R  must  finally  be  present  in  the  solution  in  order 
to  produce  the  visible  change  which  is  characteristic  of  the  indicator.  In  such 
cases  reaction  (2)  may  be  written  /i^/2,  and  the  appearance  of  the  end-point 
is  dependent  upon  the  concentration  of  72,  which  in  turn  depends  upon  that  of  R 
in  the  solution. 


COURSE  IN  QUANTITATIVE  ANALYSIS 

Ry  also,  the  least  possible  concentration  of  IR  should  cause  a 
marked  change  in  the  appearance  of  the  solution.  It  is  further 
necessary  that  the  small  amounts  of  IR  which  are  formed  locally 
during  the  titration,  in  consequence  of  imperfect  mixing,  should 
react  with  X  according  to  equation  (3),  thus  preventing  the 
appearance  of  false  end-points.  This  will  of  course  insure  the 
completion  of  the  reaction  of  X  with  R  before  the  indicator 
begins  to  be  permanently  influenced  by  R.  It  will  be  seen, 
therefore,  that  the  closeness  of  agreement  between  the  observed 
and  the  true  end-point  in  any  specific  case  will  depend  upon 
the  relative  magnitudes  of  the  three  equilibrium  constants  con- 
cerned, as  well  as  upon  the  experimental  conditions,  e.g.  the 
concentration  of  the  indicator,  the  temperature,  etc. 

If,  on  the  other  hand,  the  indicator  added  reacts  with  the 
substance  undergoing  titration,  the  reactions  concerned  must 
take  place  essentially  according  to  the  following  equations : 

(a) 
(6) 
(c)  IX+R^RX+I. 

The  appearance  of  the  end-point  is  here  dependent  upon  the 
final  completion  of  reaction  (c)  from  left  to  right.  The  accuracy 
of  the  process  in  any  given  case  will  depend  here  also  upon  the 
relative  magnitudes  of  the  three  equilibrium  constants,  and 
upon  the  experimental  conditions.  In  a  few  cases,  in  which 
IX  will  not  react  with  R,  it  is  necessary  to  use  a  special  pro- 
cedure ;  as  the  end-point  is  approached,  the  solution  under- 
going titration  is  frequently  tested,  one  drop  at  a  time  on  a 
white  test-plate,  with  a  drop  of  the  indicator  solution. 

A  factor  which  is  usually  of  great  importance  in  volumetric 
analysis  is  the  concentration  of  the  indicator  in  the  solution. 
In  those  cases  in  which  the  end-points  obtained  depend  upon 
a  change  from  one  specific  color  to  another,  the  two  colors  may 
tend  to  mask  one  another  and  give  rise  to  a  series  of  indeter- 
minate transition  tints.  It  is  then  desirable  that  the  entire 


VOLUMETRIC  ANALYSIS  103 

amount  of  indicator  present  should  be  promptly  transformed 
by  the  slightest  possible  excess  of  the  titrating  solution ;  in  such 
cases  only  a  very  slight  amount  of  the  indicator  should  be  used.1 
If,  however,  the  solution  containing  the  indicator  is  colorless, 
and  the  addition  of  an  excess  of  the  standard  solution  produces 
a  specific  color,  a  relatively  larger  amount  of  indicator  is  less 
likely  to  be  harmful ;  in  some  cases,  in  which  IR  is  a  substance 
subject  to  dissociation,  it  is  necessary,  in  order  by  mass  action 
to  insure  its  prompt  formation,  to  add  the  indicator  in  consider- 
able amount.  (Cf .  the  use  of  starch  in  iodometric  methods,  and 
also  the  use  of  ferric  alum  in  the  titration  of  silver  with  thio- 
cyanate  solutions.) 

The  Advantages  of  the  Volumetric  System.  Volumetric 
determinations  can  usually  be  carried  out' much  more  rapidly 
and  conveniently  than  the  corresponding  gravimetric  processes. 
The  actual  titration  requires  a  few  minutes  only,  but  the  neces- 
sity of  removing  interfering  substances  and  of  transforming  the 
substance  itself  into  a  form  suitable  for  titration  often  increases 
the  time  of  the  analysis  to  several  hours.  In  many  cases  vol- 
umetric processes  are  more  accurate  than  the  corresponding 
gravimetric  processes,  but  in  other  cases  the  reverse  is  true. 
Volumetric  processes  often  avoid  the  errors  which  are  involved 
in  making  gravimetric  determinations;  that  is,  the  errors  re- 
sulting from  solubility,  from  the  contamination  of  precipitates, 
and  from  actual  mechanical  losses.  On  the  other  hand  they 
necessarily  involve  certain  errors  in  the  preparation  and  measure- 
ment of  the  standard  solutions  (see,  however,  what  is  said  in 
Part  I  concerning  the  use  of  weight  burettes)  and  in  the  deter- 
mination of  the  end-points  of  the  reactions  which  are  utilized. 

General  Directions.  For  successful  volumetric  work  it  is 
essential  that  uniformity  of  practice  prevail  throughout  with 

1  In  the  case  of  methyl  orange,  for  example,  the  indicator  is  prepared  by  dis- 
solving 0.02-0.03  g.  of  the  solid  compound  in  100  cc.  of  water,  and  in  any  one 
titration  only  about  three  drops  of  this  solution  should  be  used.  Counting  20 
drops  to  i  cc.,  and  using  3  drops  in  a  titration,  the  amount  of  methyl  orange  actually 
present  is  less  than  0.05  mg. 


104     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

respect  to  all  matters  which  can  influence  the  accuracy  of  meas- 
urement of  liquids.  For  example,  whatever  time  is  allowed  for 
drainage  in  the  calibration  of  a  measuring  vessel,  should  also 
be  allowed  whenever  the  vessel  is  used ;  and  parallel  conditions 
should  be  insured  during  both  standardization  and  analysis, 
with  respect  to  the  quantity  of  the  indicator  and  the  final  volume 
of  the  reaction  mixture  after  titration.  In  some  cases  the  stand- 
ard of  the  solution  will  vary  appreciably  with  variation  of  the 
experimental  conditions. 

It  is  of  course  necessary  that  standard  solutions  should  be 
protected  from  concentration  or  dilution;  they  should  be  kept 
stoppered  and  away  from  direct  sunlight  or  heat.  The  bottles 
should  be  shaken  before  use  to  collect  any  water  that  may  have 
evaporated  from  the  solution  and  condensed  on  the  sides,  and, 
after  use,  before  replacing  the  stoppers,  the  necks  and  stoppers 
of  the  bottles  should  be  wiped  dry  with  a  clean,  lintless  towel. 

Needless  to  say,  the  measuring  vessels  must  be  clean  and  free 
from  grease,  and  great  care  must  be  taken  to  thoroughly  rinse 
out  all  burettes  and  pipettes  with  the  standard  solutions  they 
are  to  contain,  in  order  to  remove  all  traces  of  water  or  other 
liquid,  which  would  act  as  a  diluent.  It  is  best  to  rinse  them 
three  times  with  small  portions  of  the  solution,  allowing  each 
portion  to  run  out  through  the  tip,  before  assuming  them  to 
be  in  a  condition  to  be  filled  and  used. 

Much  time  may  be  saved  by  estimating,  if  possible,  before 
beginning  the  operation,  the  approximate  volume  of  standard 
solution  which  will  be  required  for  the  titration.  This  makes 
it  possible  to  run  in  rather  rapidly  almost  the  required  amount, 
after  which,  of  course,  the  end-point  must  be  very  carefully 
determined.  In  case  such  a  calculation  cannot  be  made,  it  is 
often  worth  while  to  ascertain  this  approximate  volume  by 
means  of  a  very  rapid  preliminary  titration :  this  of  course  will 
necessitate  the  use  of  an  extra  sample.  (See,  for  example,  the 
determination  of  manganese  by  titration  with  potassium  per- 
manganate.) 


VOLUMETRIC  ANALYSIS  105 

A.  NEUTRALIZATION  METHODS 

ALKALIMETRY  AND   ACIDIMETRY 

Standard  solutions  of  acid  and  alkali  are  required  for  these 
processes,  together  with  suitable  indicators. 

Standard  Acid  Solutions.  These  are  generally  prepared  from 
hydrochloric  or  sulphuric  acid.  Hydrochloric  acid  has  the 
advantage  of  forming  soluble  compounds  with  the  alkali  earths, 
but  its  solutions  cannot  be  boiled  without  danger  of  loss.  Both 
acids  may  be  used  with  all  indicators. 

Standard  Alkali  Solutions.  These  may  be  prepared  from 
sodium  or  potassium  hydroxide,  sodium  carbonate,  or  barium 
hydroxide.  Sodium  and  potassium  hydroxides,  if  free  from 
carbonate,  may  be  used  with  all  indicators,  but  they  absorb 
carbon  dioxide  readily  and  attack  the  glass  of  bottles;  sodium 
carbonate  may  be  weighed  directly  for  the  preparation  of  stand- 
ard solutions,  provided  its  purity  is  assured,  but  with  many 
indicators  the  liberation  of  carbonic  acid  is  a  disadvantage. 
Barium  hydroxide  solutions  are  free  from  carbon  dioxide;  if 
any  of  this  gas  is  absorbed,  it  at  once  causes  the  formation  of  a 
precipitate.  Barium  hydroxide  may  be  used  with  all  indicators, 
but  it  is  not  freely  soluble  in  water.  In  many  cases  in  which  it 
is  desirable  to  have  a  more  concentrated  carbonate-free  solution 
of  alkali,  a  sodium  hydroxide  solution  from  which  the  carbonate 
has  been  removed  by  means  of  a  slight  excess  of  barium  chloride 
is  very  serviceable.  Carbonate-free  solutions  should  be  pro- 
tected by  means  of  a  soda-lime  absorption  tube. 

Half-normal  or  tenth-normal  solutions  are  employed  in  most 
analyses,  the  latter  strength  being  convenient  when  small  amounts 
of  acid  or  alkali  are  to  be  determined. 

Indicators  for  Use  in  Alkalimetry  and  Acidimetry.  The  in- 
dicators used  in  these  processes  are  organic  substances,  often 
of  very  complicated  structure,  each  being  capable  of  existence 
in  two  forms  of  different  color,  which  are  mutually  transform- 
able into  one  another  at  specific  H+-  and  OH~-ion  concentrations. 


io6     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

An  aqueous  solution  which  is  neutral  contains  H+  and  OH~ 
ions  in  equal  concentration,  and,  so  long  as  the  solution  is  very 
dilute,  the  number  of  mols  of  water  ionized  in  one  liter  is 
o.ooooooi,  or  io~7,  at  25°.  Since  in  the  ionization  equilibrium, 
H2O«±H++OH-,  we  must  have  (H+)  (OH-)  =  fc,  it  follows  that 
in  dilute  solutions  (H+)(OR-)  =  io~14. 

If  an  acid  is  added  to  water,  or  to  a  dilute  neutral  solution, 
in  sufficient  amount  to  increase  the  H+-ion  concentration  from 
io~7  to  io~6,  then  the  OH~-ion  concentration  will  fall  to  io~8. 
Only  0.0000009  mol  of  H+  ion  would  be  required  to  produce 
such  a  change  in  a  liter  of  neutral  solution,  and  this  amount  is 
contained  in  about  o.oi  cc.  of  tenth-normal  hydrochloric  acid, 
while  the  addition  of  only  o.i  cc.  to  one  liter  would  increase  the 
H+-ion  concentration  from  io~7  to  io~5. 

The  color  changes  which  are  characteristic  of  these  indicators 
do  not  as  a  rule  occur  in  exactly  neutral  solutions.  The  accom- 
panying table  gives  for  some  of  the  more  common  indicators, 
the  H+-ion  concentrations  at  which  the  color  changes  occur. 

It  might  seem,  at  first  thought,  that  only  an  indicator  which 
changes  exactly  at  the  neutral  point  would  be  suitable,  but  this 
is  by  no  means  true.  In  titrating  a  strong  acid  against  a  strong 
base  a  few  hundredths  of  a  cubic  centimeter  of  tenth-normal 
acid  or  alkali  in  excess  will  carry  the  concentration  of  hydrogen 
or  hydroxide  ion  so  far  to  one  side  of  the  neutral  point  that  any 
of  the  indicators  for  which  the  characteristic  point  lies  between 
io~5  and  io~9  will  give  a  sharp  and  accurate  end-point. 

In  titrating  a  weak  acid  with  a  strong  base,  e.g.  acetic  acid 
with  sodium  hydroxide,  the  acetate  ions  from  the  highly  ionized 
sodium  acetate  drive  the  reaction,  HC2H302<^H++C2H302~, 
far  to  the  left,  long  before  an  equivalent  quantity  of  the  base 
has  been  added;  the  concentration  of  hydrogen  ion  becomes 
exceedingly  low  (<io~5  but  >io~7)  before  the  major  portion  of 
the  acetic  acid  has  had  its  hydrogen  replaced.  In  such  a  case 
the  change  in  color  for  methyl  orange  (see  the  Table)  will  appear 
gradually,  during  the  addition  of  a  cubic  centimeter  or  more  of 


VOLUMETRIC  ANALYSIS 


107 


3 


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io8     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  alkali,  long  before  an  equivalent  quantity  of  alkali  has  been 
run  in,  and  no  sharp  end-point  will  be  indicated.  If  phenol- 
phthalein  is  used,  however,  the  change  in  color  will  not  occur 
until  the  true  neutral  point,  (H+)  =  io~7,  has  been  passed ;  a  very 
slight  excess  of  alkali  will  then  reduce  the  concentration  of  the 
hydrogen  ion  far  below  io~8,  a  sharp  end-point  will  be  indicated, 
and  the  total  alkali  added  will  correspond  very  exactly  to  that 
required  for  the  actual  neutralization  of  the  acid. 

When  a  weak  base  is  titrated  with  a  strong  acid,  e.g.  ammonium 
hydroxide  with  hydrochloric  acid,  the  conditions  are  reversed 
and  such  indicators  as  methyl  orange,  methyl  red,  or  cochineal, 
which  change  color  in  a  faintly  acid  solution  (i.e.  H+>io~7)  are 
most  suitable. 

The  following  rules  should  be  observed  in  the  use  of  neutraliza- 
tion methods  for  the  determination  of  acids  and  bases : 

1.  In  the  titration  of  a  strong  acid  with  a  strong  base,  or  vice 
versa,  use  any  indicator  in  the  list,  from  methyl  orange  to  phenol- 
phthalein. 

2.  In  the  titration  of  any  acid  other  than  a  strong  mineral 
acid  with  a  strong  base,  use  phenolphthalein,  trinitrobenzene, 
or  a  similar  indicator. 

3.  In  the  titration  of  a  weak  base  with  a  strong  acid,  use 
methyl  orange,  Congo  red,  or  a  similar  indicator. 

4.  Do  not  attempt  to  titrate  a  weak  base  against  a  weak  acid. 
The  sensitivity  of  a  given  indicator  may  vary  under  widely 

differing  conditions  of  temperature  and  dilution,  and  for  that 
reason  it  is  important  to  titrate  approximately  equal  volumes  of 
solution  in  standardization  and  in  analysis;  and  when  it  is 
necessary,  as  is  often  the  case,  to  titrate  the  solution  at  the 
boiling  temperature,  the  standardization  should  take  place 
under  the  same  conditions.  It  is  also  obvious  that  since  some 
acid  or  alkali  is  required  to  produce  the  change  in  the  indi- 
cator itself,  the  amount  of  indicator  used  should  be  uniform 
and  not  excessive ;  usually  three  or  four  drops  of  the  solution 
are  ample. 


VOLUMETRIC  ANALYSIS  109 

Methyl  orange  solution  is  most  readily  prepared  by  dissolv- 
ing 0.02-0.05  g-  °f tne  s°lid  compound  (also  known  as  Orange  III) 
in  a  very  little  alcohol  and  diluting  with  water  to  100  cc.  It 
may  be  successfully  used  for  the  titration  of  strong  acids  and 
bases,  and  is  particularly  useful  in  the  determination  of  weak 
bases,  such  as  ammonium  hydroxide  and  certain  weak  organic 
bases.  It  can  also  be  used  in  titrating  with  a  strong  acid  the 
soluble  salts  of  very  weak  acids,  such  as  carbonates,  sulphides, 
arsenites,  borates,  and  silicates,  because  in  such  cases  the  acids 
which  are  liberated  are  too  weak  to  affect  the  indicator,  and  the 
reddening  of  the  solution  does  not  take  place  until  a  very  slight 
excess  of  the  strong  acid  has  been  added.  It  should  be  used  in 
cold,  not  too  dilute  solutions.  Its  sensitivity  is  less  in  the 
presence  of  large  quantities  of  alkali  salts,  or  of  smaller  quan- 
tities of  certain  other  metallic  salts. 

Phenolphthalein  solution  is  prepared  by  dissolving  one  gram 
of  the  solid  compound  in  100  cc.  of  alcohol.  This  indicator  is 
particularly  valuable  in  the  determination  of  weak  acids,  es- 
pecially of  organic  acids.  It  should  not  be  used  with  ammonia 
or  weaker  bases.  It  is  decolorized  by  carbonic  acid,  which  must 
therefore  be  removed  by  heating  when  other  substances  are 
being  determined ;  unlike  methyl  orange,  it  is  sensitive  in  boiling- 
hot  solutions  The  volume  of  the  solutions  titrated  should  be 
approximately  uniform  in  standardization  and  in  analysis,  and 
for  the  best  results  this  volume  should  not  in  general  exceed 
125-150  cc. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROXI- 
MATELY HALF-NORMAL  SOLUTIONS  OF  HYDROCHLORIC 
ACID  AND  SODIUM  HYDROXIDE 

Procedure.  Pour  into  a  zoo-cc.  measuring  cylinder  a  volume 
of  hydrochloric  acid  (sp.  gr.  1.19,  with  37.2%  of  HC1)  sufficient 
to  contain  36.5  g.  of  hydrogen  chloride,  transfer  it  quantitatively 
to  a  one-liter  measuring  flask,  and  dilute  to  the  mark  with  dis- 
tilled water.  Pour  this  solution  without  loss  into  a  clean,  well- 


no     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

drained  2§  liter  bottle,  refill  the  measuring  flask  with  distilled 
water,  and  pour  this  also  into  the  bottle.  Finally  shake  the 
solution  thoroughly  for  a  full  minute  in  the  stoppered  bottle, 
to  insure  uniformity  of  concentration. 

Weigh  out,  upon  a  rough  balance,  about  42  g.  of  stick  sodium 
hydroxide.  Dissolve  the  alkali  in  water,  in  a  large  beaker, 
dilute  this  solution  also  to  two  liters,  and  shake  for  a  full  minute 
in  the  stoppered  bottle. 

Fill  two  clean,  grease-free  burettes  with  the  respective  solu- 
tions, first  thoroughly  rinsing  each  burette  three  times  with 
10  cc.  portions  of  the  corresponding  solution  and  allowing  the 
wash  liquid  each  time  to  run  out  through  the  tip.  See  that  all 
air  bubbles  are  expelled  from  the  tips,  note  the  exact  position  of 
the  liquid  in  each  burette,  and  record  the  readings  in  the  note- 
book. Run  out  from  one  burette  about  20  cc.  of  the  acid  into 
a  beaker,  and  add  three  drops  of  methyl  orange  solution ;  dilute 
the  acid  to  about  80  cc.,  and  run  in  alkali  solution  from  the 
other  burette,  with  stirring,  until  the  color  of  the  solution  changes 
from  pink  to  yellow.  Wash  down  the  sides  of  the  beaker  with 
a  little  distilled  water,  replace  the  beaker  under  the  acid  burette, 
and  add  acid  to  restore  the  pink;  continue  these  alternations 
until  the  point  is  accurately  fixed  at  which  a  single  drop  of  either 
solution  will  produce  a  distinct  change  of  color.  Select  as  the 
end-point  either  the  appearance  of  the  faintest  tinge  of  pink 
which  can  be  recognized,  or  its  disappearance,  and  always  titrate 
to  the  same  point.  If  the  titration  has  occupied  more  than  three 
minutes,  the  time  required  for  draining,  the  readings  of  the 
burettes  may  be  taken  at  once  and  entered  in  the  notebook. 

Refill  the  burettes  and  repeat  the  titration.  Correct  the 
burette  readings  as  indicated  by  the  burette  calibrations,  and  if 
necessary  for  temperature  (see  Part  I),  and  obtain  the  ratio  of 
the  solutions  as  shown  in  the  following  example : 

cc.  acid      21.53  r      -j  r    n    T 

—  =  — —  =  1.022  cc.  of  acid  per  i.ooo  cc.  of  alkali, 
cc.  alkali     21.07 


VOLUMETRIC  ANALYSIS  in 

When  this  ratio  has  been  satisfactorily  established,  the  hydro- 
chloric acid  solution  is  standardized  as  follows  : 

(a)  By  Titration  against  Pure  Sodium  Carbonate.    If  a  suit- 
able oven  is  available,  dry  the  salt  on  a  watch  glass  for  an  hour 
at  130-150° ;  otherwise  heat  about  5  g.  of  pure  sodium  carbonate 
in  a  small  porcelain  dish,  on  a  wire  gauze  over  a  small  Bunsen 
flame,  for  one  half  hour,  and  then  allow  the  salt  to  cool  in  a 
desiccator.     Transfer   the  cold   salt  to  a  dry,   well-stoppered 
weighing  tube,  and  weigh  out  into  4oo-cc.  beakers  two  portions 
of  0.5-0.6  g.  each,1  noting  the  exact  weights  ,in  the  notebook. 
Pour  over  the  salt  about  80  cc.  of  water,  stir  until  dissolved,  and 
add  three  drops  of  methyl  orange  solution.     Fill  the  burettes 
with  the  acid  and  alkali  solutions,  note  the  initial  readings  of 
the  burettes,  and  run  in  the  acid,  with  stirring,  until  the  solution 
assumes  the  faintest  tinge  of  pink.     Wash  down  the  sides  of  the 
beaker  with  a  little  water,  and  if  the  solution  loses  its  pink  tint 
add  acid,  one  drop  at  a  time,  with  stirring,  until  the  faint  pinkish 
tinge  just  returns.     (If  too  much  acid  is  added,  the  excess  can 
be  determined  by  means  of  the  alkali  in  the  other  burette,  since 
the  ratio  of  the  alkali  to  the  acid  is  known.)     After  three  minutes, 
note  the  burette  readings  and  enter  them  in  the  notebook. 
From  the  data  recorded,  calculate  the  normality  factor  of  the 
acid  and  also  that  of  the  alkali.     The  standardization  of  the 
acid  must  be  repeated  until  the  duplicate  values  agree  within 
two  parts  in  one  thousand,  and  the  same  of  course  applies  to 
the  determination  of  the  ratio  of  the  two  solutions. 

(b)  Gravimetrically  with  Silver  Nitrate.     Measure   out  accu- 
rately from  a  pipette  10.00  cc.  of  the  acid  into  each  of  two 
3OO-CC.  beakers,  and  dilute  in  each  case  with  150  cc.  of  water. 
Precipitate  the  chlorine  from  these  solutions  with  silver  nitrate 
according  to  the  procedure  given  in  Part  II,  and  filter  the  silver 
chloride  off  through  Gooch  crucibles,  prepared  and  weighed  as 
there  indicated.     Wash  the  precipitates  with  hot  water  until 

1  The  weights  of  samples  in  this  book  are  based  upon  the  use  of  30-00.  burettes. 
If  50-cc.  burettes  are  used,  it  is  better  to  take  samples  f  as  large. 


112     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

free  from  soluble  silver  salts,  dry  at  120-130°  to  constant  weight, 
and  from  the  weight  of  silver  chloride  found  in  each  case  calcu- 
late the  normality  factor  of  the  acid.  The  duplicate  values 
should  agree  very  closely  with  one  another,  and  also  with  those 
previously  found  by  titration  against  sodium  carbonate. 

NOTES.  —  i.  Although  silver  chloride  is  insoluble,  the  normality  factor 
of  the  hydrochloric  acid  may  nevertheless  be  calculated  directly  from  the 
weight  of  the  precipitate  obtained.  For  example,  if  10.00  cc.  of  the  acid 
were  found  to  yield  0.7317  g.  of  AgCl,  then  (since  an  equal  volume  of  normal 
HC1  would  yield  1.4334  g.  of  AgCl)  the  normality  factor  of  the  acid  is 
0.7317/1.4334,  or  0.5105. 

In  the  same  way,  the  normality  factor  of  the  acid  may  be  calculated  in 
the  case  of  Method  (a).  For  example,  if  0.5682  g.  of  Na2C03  require 
21.00  cc.  of  acid,  the  normality  factor  is  equal  to  0.5682/1.1130=0.5105 
(1.1130  g.  is  the  amount  of  Na2C03  contained  in  21.00  cc.  of  the  normal 
solution). 

If  it  has  previously  been  found,  for  example,  that  i.ooo  cc.  of  alkali 
solution  is  equivalent  to  1.022  cc.  of  the  acid,  then  it  follows  that  the  nor- 
mality factor  of  the  alkali  is  0.5105X1.022  =  0.5217. 

2.  If  it  is  desired  to  prepare  solutions  of  exactly  one  half  normal  con- 
centration, slightly  stronger  solutions  are  first  prepared,  and,  after  stand- 
ardization, they  are  diluted  with  the  calculated  volume  of  water.     For 
example,  the  0.5105  N  acid  should  be  diluted  according  to  the  proportion, 
0.5105 : 0.5000= x :  icoo,  and  the  0.5217  N  alkali  according  to  the  proportion, 
0.5217:  0.5000=37 :  1000;  i.e.  one  liter  of  each  solution  should  have  added 
to  it  21.0  cc.  and  43.2  cc.  of  water,  respectively.    The  water  is  added  from 
a  burette.    After  dilution  the  solutions  should  be  thoroughly  shaken,  and 
then  restandardized. 

3.  Solutions  should  be  thoroughly  mixed  to  insure  uniformity  of  con- 
centration before  standardization.    They  should  be  allowed  to  attain  the 
temperature  of  the  room,  and  they  should  be  shaken  to  take  up  any  water 
which  may  have  evaporated  and  later  condensed  on  the  inner  walls  of  the 
bottles.    Before  replacing  the  stopper  in  a  bottle,  always  dry  it,  as  well 
as  the  neck  of  the  bottle,  with  a  clean,  lintless  towel. 

4.  The  liquid  is  diluted  to  100  cc.  during  standardization  in  order  that 
the  volume  may  be  the  same  as  that  which  will  prevail  during  analysis. 

5.  The  exact  point  at  which  the  color  changes  should  be  chosen  as  the 
end-point;  any  deeper  tint  is  unsatisfactory,  since  it  is  not  possible  to 
duplicate  shades  of  color  from  day  to  day  by  memory. 


VOLUMETRIC  ANALYSIS  113 

6.  The  selection  of  the  best  compound  to  be  used  as  a  standard  for  acid 
solutions  has  been  the  subject  of  much  controversy.    In  many  works  cal- 
cium carbonate  has  been  recommended  (see  Part  IV,  Problem  54),   and 
also  sodium  carbonate,  which  can  now  be  purchased  sufficiently  pure  for 
this  purpose.    The  most    reliable   standard   would   seem   to   be   sodium 
carbonate  prepared  from  recrystallized  sodium  bicarbonate  by  heating 
the  latter  between  280°  and  300°.    The  bicarbonate  is  easily  purified  by 
crystallization,  and  between  the  temperatures  named  it   is  decomposed 
quantitatively  according  to  the  equation, 

2  NaHCO3=Na2C03+H20+C02. 

7.  Instead  of  standardizing  the  acid  solution,  it  is  equally  practicable 
to  standardize  the  alkali  solution,  with  the  use  of  phenolphthalein,  against 
any  of  the  following  pure  acids :  oxalic  acid,  H2C2O4 .  2  H2O ;  acid  potas- 
sium oxalate,  KHC204 .  H20 ;  potassium  tetroxalate,  KH3(C204)2 .  2  H2O ; 
potassium  bitartrate,  KHC4H406 ;  succinic  acid,  H2C4H404.     The  last  two 
are  probably  the  most  suitable,  since  they  are  free  from  water  of  crystalliza- 
tion.   It  should  be  noted  that  the  acid  oxalate  and  the  bitartrate  contain 
one  replaceable  hydrogen  atom  each,  while  the  tetroxalate  contains  three 
such  atoms,  and  the  oxalic  and  succinic  acids  two. 

8.  While  it  is  permissible  to  standardize  hydrochloric  acid  solutions 
(provided  they  are  free  from  other  chlorides)  gravimetrically  with  silver 
nitrate,  a  solution  of  sulphuric  acid  should  not  be  standardized  in  this  way 
by  precipitation  as  barium  sulphate ;  the  results  would  be  less  reliable  on 
account  of  the  difficulty  in  obtaining  large  precipitates  of  barium  sulphate 
which  are  free  from  contamination  and  which  are  not  partially  reduced 
on  ignition. 

THE  DETERMINATION  OF  THE  TOTAL  ALKALINE 
VALUE  OF   SODA  ASH 

The  sample  may  be  one  of  commercial  soda  ash,  or  it  may  be 
an  artificial  mixture  of  sodium  carbonate  and  sodium  chloride. 

Procedure.  Weigh  out  roughly  on  a  watch  glass  5  g.  of  the 
soda  ash,  and  dry  it  for  one  hour  at  110° ;  allow  it  to  cool  in  a 
desiccator.  Now  accurately  weigh  the  sample  on  the  watch 
glass,  transfer  it  quantitatively  to  a  5oo-cc.  beaker,  washing  off 
the  glass  with  about  50  cc.  of  water ;  dry  and  weigh  the  watch 
glass,  and  take  the  difference  as  the  weight  of  the  sample.  Gently 
warm  the  sample  with  the  50  cc.  of  water  and  filter  off  any  in- 


114     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

soluble  residue.  Wash  the  filter  at  least  five  times  with  20  cc. 
portions  of  warm  water,  and  receive  the  filtrate  and  washings 
in  a  250-00.  measuring  flask  which  has  been  freed  from  grease 
by  means  of  cold  cleaning  solution.  Cool  the  liquid  to  the  room 
temperature,  add  distilled  water  until  the  lowest  point  of  the 
meniscus  is  level  with  the  mark  on  the  neck  of  the  flask,  and 
thoroughly  mix  the  solution  by  pouring  it  from  the  flask  into  a 
dry  beaker  and  back  into  the  flask  two  or  three  times. 

Measure  off  25  cc.  of  the  solution  with  a  pipette  (which  should  be 
previously  rinsed  out  with  small  quantities  of  the  solution)  into  a 
300-cc.  Erlenmeyer  flask,  allowing  the  pipette  to  drain  for  a  sec- 
ond or  two  with  the  tip  in  contact  with  the  inside  wet  surface  of 
the  flask  (unless  it  was  standardized  otherwise).  Dilute  the  solu- 
tion to  about  80  cc.,  add  three  drops  of  methyl  orange  solution, 
and  titrate  with  the  standard  acid,  using  the  standard  alkali 
to  complete  the  titration  as  already  described.  From  the 
volumes  of  acid  and  alkali  used,  corrected  for  temperature  dif- 
ference and  burette  errors,  calculate  the  percentage  of  alkali 
present,  assuming  the  alkali  to  be  wholly  sodium  carbonate. 

Measure  out  other  portions  of  25  cc.  from  the  main  solution, 
and  repeat  the  titration  until  satisfactory  checks  are  obtained. 

NOTES.  —  i.  Let  us  assume,  for  example,  that  5.890  g.  of  soda  ash  were 
used  in  the  preparation  of  250.0  cc.  of  solution,  and  that  25.75  cc.  of  0.5105  N 
acid  and  4.13  cc.  of  0.5217  N  alkali  were  used  in  the  titration  of  25.00  cc.  of 
this  solution.  Then  it  follows  that  5.890  g.  of  the  soda  ash  are  equivalent 
to  ioX(25. 75X0.5105— 4.13X0.5217)  =  109.90  cc.  of  N  acid;  and,  since 
this  volume  of  normal  acid  would  neutralize  an  equal  volume  of  normal 
alkali,  therefore  the  5.890  g.  of  soda  ash  contained  0.053 X  109.90=  5.8246  g., 
or  98.9%  of  Na2C03. 

2.  Soda  ash  is  crude  sodium  carbonate.    When  made  by  the  Solvay 
process  it  is  apt  to  contain  also  sodium  chloride,  sulphate,  and  either  the 
bicarbonate  or  hydroxide ;  if  made  by  the  Le  Blanc  process,  which  however 
has  gone  out  of  use,  sodium  sulphide,  silicate,  aluminate,  and  other  impurities 
are  likely  to  be  present.     Many  of  these  contribute  to  the  total  alkaline  value? 
but  it  is  customary  to  calculate  this  value  in  terms  of  sodium  carbonate  alone. 

3.  In  order  to  obtain  uniform  results,  it  is  customary  to  dry  the  soda 
ash  at  110°  before  analysis.     Complete  expulsion  of  the  moisture  would 


VOLUMETRIC  ANALYSIS  115 

require  a  very  much  higher  temperature.  At  least  5  g.  are  taken,  in  order 
to  secure  a  representative  sample ;  but  since  this  is  too  much  for  convenient 
titration,  an  aliquot  portion  of  the  solution  is  measured  off. 

4.  For  other  methods  of  analyzing  soda  ash,  the  student  should  refer 
to  Part  IV,  Problems  60,  61,  and  95. 

THE  DETERMINATION  OF  THE  NEUTRALIZATION 
VALUE  OF  AN  ACID 

Procedure.  Weigh  out  accurately  into  3oo-cc.  beakers  two 
0.6-0.7  g.  portions  of  the  unknown  acid  (oxalic  acid,  acid  potas- 
sium oxalate,  potassium  tetroxalate,  potassium  bitartrate, 
succinic  acid,  or  some  similar  compound),  dissolve  each  sample 
in  about  80  cc.  of  warm  water,  add  two  or  three  drops  of  phenol- 
phthalein  solution,  and  run  in  half-normal  alkali  from  a  burette 
until  the  solution  is  pink.  Add  half-normal  acid  from  the  other 
burette  until  the  pink  color  just  disappears,  and  then  exactly 
0.30  cc.  in  excess.  Heat  the  solution,  and  boil  for  three  minutes. 
If  the  pink  color  reappears  upon  boiling,  discharge  it  with  acid, 
again  add  0.30  cc.  in  excess,  and  repeat  the  boiling.  Discharge 
the  pink  color  if  it  again  reappears,  again  adding  0.30  cc.  in 
excess.  Repeat  this  treatment  until  the  pink  color  fails  to  return 
upon  boiling  for  three  minutes.  Finally  add  alkali  until  the 
color  just  reappears,  then  a  drop  or  two  of  acid  in  excess  and  boil 
for  one  minute.  If  no  color  appears  during  this  time,  complete 
the  titration  with  alkali  in  the  hot  solution.  From  the  corrected 
volume  of  alkali  required  to  react  with  the  acid  solution,  calculate 
the  normality  factor  of  a  solution  containing  10.00  g.  of  the  un- 
known acid  in  one  liter.  The  results  should  check  within  two 
parts  in  one  thousand. 

NOTES.  —  i.  Although  it  is  desirable  to  employ  the  same  indicator 
throughout  standardization  and  analysis,  the  difference  resulting  from  the 
change  of  indicator  is  in  this  instance  insignificant,  and  the  student  may 
neglect  it.  It  should  be  remembered,  however,  that  in  order  to  obtain  the 
greatest  accuracy  possible,  a  restandardization  throughout  would  be  essential. 

2.  Since  commercial  sodium  hydroxide  always  contains  some  carbonate, 
and  since  phenolphthalein  is  sensitive  to  carbonic  acid,  the  solution  must 
i 


n6     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

be  boiled  to  free  it  from  this  acid.  Phenolphthalein  does  not  show  an 
alkaline  reaction  with  cold  dilute  sodium  bicarbonate  solution ;  hence  cold, 
dilute  solutions  of  sodium  carbonate  become  colorless  with  this  indicator 
as  soon  as  the  carbonate  has  been  transformed  into  bicarbonate  by  the 
acid.  Upon  boiling,  the  bicarbonate  is  partially  hydrolyzed  according  to 
the  equation,  Na+HCO3-+H+OH-5:Na++C>H-+H2C03,  the  solution 
loses  carbon  dioxide,  and  the  pink  color  returns.  This  must  again  be  dis- 
charged, the  solution  boiled,  and  so  on. 

3.  Hydrochloric  acid  is  volatilized  from  aqueous  solutions  upon  boiling, 
unless  they  are  very  dilute.    If  the  excess  of  acid  added  is  not  more  than 
0.30  cc.,  however,  no  loss  need  be  feared. 

4.  When  a  large  number  of  acidimetric  determinations  are  to  be  made 
with  phenolphthalein  as  the  indicator,  it  is  well  worth  while  to  prepare  and 
standardize  a  carbonate-free  alkali  solution.    The  acid  solutions  are  in 
such  cases  boiled,  to  free  them  from  carbonic  acid,  or  they  are  made  up  with 
freshly  boiled  water,  and  titrated  hot  with  carbonate-free  alkali  (cf.  the 
general  discussion  under  alkalimetry  and  acidimetry). 


THE  DETERMINATION  OF  PROTEIN  NITROGEN  BY  THE 
KJELDAHL  METHOD 

Principle.  When  an  organic  substance  is  heated  with  con- 
centrated sulphuric  acid,  especially  in  the  presence  of  an  oxygen 
carrier,  the  organic  substance  is  completely  decomposed,  and 
any  protein  (or  other  similarly  combined)  nitrogen  is  converted 
into  ammonia.  This  at  once  combines  with  acid  to  form  am- 
monium acid  sulphate,  NH4HS04,  which  remains  in  solution 
in  the  sulphuric  acid.  Upon  diluting  the  mixture  with  water 
and  adding  sodium  hydroxide  in  excess,  the  ammonia  is  liberated, 
and  can  be  distilled  over  and  collected  in  a  known  volume  of 
standard  acid,  which  it  partially  neutralizes.  By  titrating  the 
excess  of  acid  with  a  standard  alkali,  the  volume  of  the  standard 
acid  neutralized  by  the  ammonia  can  be  found,  and  from  the 
data  obtained  the  percentage  of  nitrogen  in  the  sample  may  be 
calculated. 

Procedure.  Accurately  weigh  out  from  a  weighing  tube, 
upon  separate  sheets  of  quantitative  filter  paper,  two  samples 
of  about  i  g.  each  of  the  substance  to  be  analyzed.  Wrap  each 


VOLUMETRIC  ANALYSIS  117 

sample  carefully  in  the  paper,  and  introduce  the  bundle  into  a 
clean  500  cc.  Kjeldahl  flask.  To  each  flask  add  about  0.5  g.  of 
powdered  copper  sulphate,  and  25  cc.  of  concentrated  sulphuric 
acid.  See  that  the  samples  are  thoroughly  wet  by  the  acid,  and 
then  place  the  flasks  on  the  digestion  rack  in  the  Nitrogen  Lab- 
oratory, with  the  necks. resting  in  the  circular  openings  of  the 
lead  ventilating  pipe;  place  the  flasks  in  unoccupied  positions 
as  near  as  possible  to  one  of  the  exhaust  flues.  Heat  gently 
until  frothing  ceases,  add  10  g.  (weighed  roughly)  of  potassium 
sulphate,  or  an  equivalent  weight  of  sodium  sulphate,  and 
heat  to  gentle  ebullition  for  two  or  three  hours  until  the  liquid 
is  of  a  clear  green  color,  without  any  trace  of  brown  (do  not 
allow  the  flame  to  reach  above  the  surface  of  the  liquid).  Con- 
tinue the  heating  for  half  an  hour  longer,  and  allow  to  cool. 

While  the  flasks  are  cooling,  accurately  measure  from  a  burette 
two  30.00-cc.  portions  of  0.5  N.  hydrochloric  acid,  into  4oo-cc. 
Erlenmeyer  flasks,  and  add  to  each  about  25  cc.  of  distilled  water. 
Place  these  flasks  under  the  distilling  apparatus,  so  that  the 
delivery  tubes  just  dip  into  the  acid  solutions. 

After  cooling,  carefully  dilute  the  contents  of  the  digestion 
flasks  with  150  cc.  of  distilled  water,  and  cool  again.  Carefully 
pour  down  the  inclined  neck  of  each  flask,  so  that  it  shall  not 
mix  with  the  acid  solution,  75  cc.  of  sodium  hydroxide  solution 
(300  g.  of  NaOH  per  liter).  Place  the  flasks  on  the  distilling 
rack,  add  one  or  two  pieces  of  granulated  zinc,  and  quickly 
connect  the  flasks  with  the  distilling  heads,  using  well-fitting 
rubber  stoppers.  Finally,  mix  the  contents  of  each  flask  by 
gently  rotating  it,  and  then  begin  to  heat  the  mixture. 

Distill  off  about  two  thirds  of  the  contents  of  each  flask,  with 
great  care  that  they  do  not  boil  over.  The  distillation  will  re- 
quire about  45  minutes.  Disconnect  the  distilling  flasks  and 
rinse  out  the  delivery  tubes  into  the  receiving  flasks  with  a 
little  distilled  water.  Add  3  drops  of  methyl  orange  to  each  of 
the  receiving  flasks,  and  titrate  the  contents  with  0.5  N.  sodium, 
hydroxide. 


n8     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

From  the  data  obtained,  calculate  the  percentage  of  nitrogen 
in  the  sample. 

NOTES.  —  i.  The  sulphuric  acid  hydrolyzes  the  NH2-group,  to  give 
ammonia,  and  also  acts  as  an  oxidizing  agent,  converting  the  organic  matter 
into  carbon  dioxide,  water,  or  other  volatile  products.  For  example, 
CO(NH2)2+H20+2  H2S04=C02+2  NH4HS04;  C6H1005+w  H2S04=6  C 
+5  H2O+w  H2S04;  and  C+2  H2S04=CO2+2  H20+2  S02. 

2.  The  CuS04  gives  up  oxygen  more  readily  to  the  organic  matter  than 
the  H2S04  does ;  but  the  H2S04  then  reoxidizes  the  copper  so  that  at  the 
end  of  the  operation  the  copper  is  still  present  as  copper  sulphate.    That 
is  to  say,  the  copper  salt  acts  catalytically  as  an  oxygen  carrier. 

3.  Mercuric  sulphate  is  often  used  instead  of  copper  sulphate  as  an 
oxygen  carrier,  a  few  small  globules  of  metallic  mercury  being  added  to 
the  acid  digestion  mixture.    Although  the  mercury  salt  is  somewhat  more 
efficient,  it  tenaciously  retains  ammonia,  as  H2N— Hg— 0— S02— 0— Hg 
— NH2,  even  in  the  presence  of  an  excess  of  hot  alkali,  and  it  is  therefore 
necessary  to  add  also  a  large  excess  of  sodium  sulphide.    This  converts  all 
the  mercury  into  HgS,  which  combines  with  Na2S  to  form  soluble  Hg(SNa)2, 
and  the  ammonia  is  liberated. 

4.  The  K2SO4  forms  with  the  acid  KHSO4,  and  this  serves  to  raise  the 
boiling  point  of  the  sulphuric  acid ;  the  higher  temperature  hastens  the  di- 
gestion. 

5.  The  flask  is  provided  with  a  long  neck  in  order  that  the  acid  fumes, 
which  would  otherwise  be  lost,  may  condense  and  run  back  into  the  digestion 
mixture. 

6.  After  the  acid  solution  has  been  diluted,  it  is  specifically  lighter  than 
the  NaOH  solution  used ;  upon  pouring  the  latter  carefully  down  the  neck 
of  the  inclined  flask,  it  sinks  to  the  bottom  and  leaves  the  surface  of  the 
liquid  still  acid,  thus  preventing  the  loss  of  ammonia  at  this  stage  of  the 
procedure.    The  contents  should  not  be  mixed  until  after  the  flask  has 
been  tightly  connected  with  the  distilling  head. 

7.  The  granulated  zinc  is  added  in  order  to  prevent  bumping  during 
the  distillation ;  the  zinc  dissolves  slowly  in  the  alkaline  solution,  with  the 
evolution  of  hydrogen.     Fragments  of  pumice  stone  or  of  platinum  are 
often  used  for  the  same  purpose. 

8.  If  nitrates  are  present  in  the  sample  (e.g.  a  fertilizer),  and  it  is  desired 
to  determine  the  total  nitrogen,  the  procedure  may  be  modified  as  follows : 
Thoroughly  wet  the  sample  in  the  flask  with  25  cc.  of  concentrated  sul- 
phuric acid,  in  which  one  gram  of  salicylic  acid  (CeH^^QQjj  J  has  previ- 


VOLUMETRIC  ANALYSIS  119 

ously  been  dissolved ;  this  reacts  with  the  nitric  acid  to  form  nitrosalicylic 

/N02 
acid,  CeHs — OH       .     Next  add  slowly,  with  frequent  shaking,  10  g.  of 

\COOH 

powdered  sodium  thiosulphate,  which  reduces  the  nitrosalicylic  acid  to 

/NH2 
aminosalicylic  acid,  CeHs — OH       .    Now  add  0.5  g.  of  powdered  copper 

\COOH 

sulphate  and  complete  the  determination  as  already  described,  but 
omitting  the  addition  of  the  alkali  sulphate  (the  solution  contains  NaHS04 
from  the  thiosulphate). 

9.  It  is  evident  that  ammonium  salts  may  be  analyzed  for  ammonia  by 
simply  distilling  them  with  an  excess  of  alkali,  absorbing  the  ammonia  in 
an  excess  of  standard  acid,  etc.  Moreover,  nitric  acid  and  nitrates  may 
be  quantitatively  reduced  to  ammonia  and  determined  in  this  way  (see 
Part  IV,  Problems  63,  64,  and  89). 

Certain  other  salts,  as  acetates,  may  be  analyzed  in  an  analogous  manner 
by  distillation  with  phosphoric  acid  in  excess,  the  distillate  being  collected 
in  standard  alkali  and  the  excess  of  the  latter  titrated,  with  the  use  of  phenol- 
phthalein. 

B.  METHODS  OF  OXIDATION  AND  REDUCTION 

In  most  oxidation  and  reduction  processes  the  standard  solu- 
tion employed  in  the  titration  is  one  of  an  oxidizing  agent,  though 
it  is  often  an  advantage  to  have  at  hand  a  standard  reducing 
solution  as  well.  It  may  be  stated,  in  general,  that  oxidizable 
substances  are  most  often  determined  by  direct  titration,  while 
oxidizing  substances  are  very  frequently  determined  by  indirect 
titration  (i.e.  by  adding  the  substance  to  an  excess  of  reducing 
solution,  and  titrating  either  the  excess  of  this  solution  or  a 
product,  as  iodine,  set  free  by  the  oxidizing  substance). 

Standard  Solutions.  The  most  important  oxidizing  agents 
employed  in  the  form  of  standard  solutions  are  potassium  per- 
manganate, potassium  dichromate,  iodine,  potassium  bromate, 
and  ferric  chloride;  while  the  most  important  reducing  agents 
are  ferrous  ammonium  sulphate,  oxalic  acid,  sodium  thiosul- 
phate, arsenious  acid,  and  titanous  chloride.  Other  oxidizing 
and  reducing  agents  are  frequently  used  in  the  processes,  but 
not  usually  in  the  form  of  standard  solutions. 


120     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  most  important  combinations  among  the  foregoing  stand- 
ard solutions  are:  potassium  permanganate  and  ferrous  salts 
or  oxalic  acid ;  iodine  and  sodium  thiosulphate  or  arsenious 
acid ;  potassium  dichromate  and  ferrous  salts. 

Indicators.  With  respect  to  the  indicators  employed,  potas- 
sium permanganate,  owing  to  its  own  intense  coloring  power, 
is  its  own  indicator ;  the  slightest  excess  is  indicated  with  great 
accuracy  in  otherwise  colorless  (or  even  in  certain  faintly  colored) 
reaction  mixtures.  Since  in  the  case  of  potassium  dichromate 
no  indicator  has  been  found  which  is  entirely  satisfactory  for 
use  within  the  solution,  potassium  ferricyanide  is  employed  as 
an  outside  indicator  to  determine  the  point  at  which  the  ferrous 
iron  is  completely  oxidized.  In  the  case  of  iodine,  starch  solu- 
tion is  employed  as  an  indicator.  The  use  of  these  indicators 
will  be  discussed  under  the  respective  processes. 

I.    DICHROMATE    PROCESSES 

Fundamental  Principles.  In  the  presence  of  hydrochloric 
or  sulphuric  acid,  ferrous  salts  are  promptly  and  completely 
oxidized  in  the  cold  to  ferric  salts  upon  the  addition  of  potassium 
dichromate  solution.  Since  hydrochloric  acid  is  by  far  the  most 
suitable  solvent  for  iron  and  its  compounds,  the  titration  is 
most  often  carried  out  in  the  presence  of  this  acid : 
6  FeCl2+K2Cr2O7+i4  HC1  =  6  FeCl3  +  2  KC1+2  CrCl3  +  7  H20. 

As  an  indicator,  potassium  ferricyanide  is  used  outside  the 
solution  to  determine  the  end-point  of  the  reaction.  A  drop  of 
the  iron  solution  is  added  to  one  of  the  indicator  solution  on  a 
white  surface,  and  the  mixture  examined  for  a  blue  coloration 
due  to  the  formation  of  insoluble  ferrous  ferricyanide.  The 
potassium  ferricyanide  must  of  course  be  free  from  ferrocyanide, 
and  the  indicator  solution  must  be  very  dilute  to  diminish  the 
interference  of  its  own  color ;  a  crystal  the  size  of  a  pin- 
head  dissolved  in  25  cc.  of  water  gives  the  right  concentration. 
Since  this  solution  is  not  stable,  it  must  be  freshly  prepared 
each  day. 


VOLUMETRIC  ANALYSIS  121 

THE  PREPARATION  AND  STANDARDIZATION  OF  THE  AP- 
PROXIMATELY TENTH-NORMAL  DICHROMATE  AND 
FERROUS  IRON  SOLUTIONS 

Procedure.  Pulverize  about  3  g.  of  potassium  dichromate, 
dissolve  2.5  g.  of  the  powder  in  water,  and  dilute  to  500  cc. ; 
also  dissolve  20  g.  of  ferrous  ammonium  sulphate  and  5  g.  of 
ammonium  sulphate  in  water,  with  the  addition  of  5  cc.  of  con- 
centrated sulphuric  acid,  and  dilute  to  500  cc.  Thoroughly 
mix  the  solutions,  see  that  they  are  of  the  room  temperature, 
and  then  fill  a  burette  with  each  solution,  observing  the  pre- 
cautions previously  emphasized. 

Prepare  a  solution  of  potassium  ferricyanide  of  the  strength 
recommended  above,  and  place  single  drops  of  this  solution  on 
the  surface  of  a  white  porcelain  tile.  Run  out  from  a  burette 
into  a  300-cc.  beaker  about  20  cc.  of  the  ferrous  solution,  add 
15-20  cc.  of  6-normal  hydrochloric  acid,  dilute  to  150  cc.,  and 
run  in  about  18  cc.  of  the  dichromate  solution  from  a  second 
burette.  Test  at  this  point  by  adding  a  small  drop  of  the 
well-mixed  iron  solution  from  the  end  of  a  stirring  rod  to  a 
drop  of  indicator  on  the  tile.  (The  stirring  rod  which  has 
touched  the  indicator  should  be  washed  off  with  distilled  water 
before  being  returned  to  the  iron  solution.)  If  a  blue  pre- 
cipitate appears  at  once,  0.5  cc.  of  the  dichromate  solution  may 
be  added  before  another  test  is  made.  As  soon  as  the  blue 
appears  to  be  less  pronounced,  add  the  dichromate  solution 
in  smaller  amounts,  finally  a  drop  at  a  time,  until  the  point 
is  reached  at  which  a  bluish  coloration  fails  to  appear  within 
30  seconds  after  mixing  a  large  drop  of  the  iron  solution  with 
a  drop  of  the  indicator  on  the  tile,  the  time  being  carefully 
noted.  As  soon  as  the  30  seconds  have  elapsed,  remove  another 
large  drop  of  the  iron  solution  and  mix  it  with  the  indicator 
beside  the  last  test;  if  no  difference  can  be  noted  between 
the  last  mixture  and  this  fresh  one,  the  reaction  is  com- 
plete. Should  the  end-point  accidentally  be  overstepped,  more 
of  the  ferrous  solution  may  be  added  and  the  titration  pro- 


122     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

ceeded  with  as  before.  Repeat  this  titration  until  good  dupli- 
cates are  obtained. 

From  the  volumes  of  the  solutions  used  (after  applying  any 
necessary  corrections)  calculate  the  value  of  the  ferrous  solution 
in  terms  of  the  dichromate  solution;  the  ratios  found  should 
not  differ  by  more  than  two  parts  in  one  thousand. 

Standardize  the  dichromate  solution  as  follows:  Weigh  out 
two  portions  of  bright  iron  wire  of  about  0.15  g.  each.  The 
wire  should  be  free  from  rust,  and  should  be  handled  with  filter 
paper.  It  should  be  bent  so  as  not  to  interfere  with  the  move- 
ment of  the  balance.  Place  2O-cc.  portions  of  6-normal  hydro- 
chloric acid  in  3oo-cc.  beakers,  cover  them  with  watch  glasses, 
and  heat  to  boiling.  Remove  the  flames,  drop  in  the  portions  of 
wire,  and  after  the  solution  of  the  iron  boil  carefully  for  two  or 
three  minutes,  keeping  the  beakers  covered.  Wash  the  sides 
of  the  beakers  and  the  watch  glasses  with  a  very  little  water, 
and  add  stannous  chloride  solution  to  the  hot  liquid,  drop  by 
drop,  until  the  mixture  is  colorless ;  avoid  more  than  a  drop  or 
two  in  excess.  Allow  to  cool,  dilute  with  150  cc.  of  cold  water, 
and  add  rapidly  with  stirring  25  cc.  of  mercuric  chloride  solu- 
tion. Allow  the  solutions  to  stand  for  two  minutes,  and  titrate 
without  further  delay.  (Calculate  the  volume  of  o.i  .N  dichro- 
mate solution  which  would  be  required  by  the  sample  of  wire, 
and  add  almost  this  quantity  before  beginning  to  test  with  the 
indicator.)  The  ferrous  solution  may  be  used  if  the  end-point 
is  passed. 

From  the  volume  of  dichromate  solution  required  to  oxidize 
the  known  quantity  of  iron,  calculate  the  normality  factor  of 
the  solution.  Repeat  the  standardization  until  duplicates  are 
obtained  which  do  not  differ  by  more  than  two  parts  in  a  thou- 
sand, and  from  the  mean  of  these  calculate  the  normality  factor 
of  the  ferrous  ammonium  sulphate  solution. 

NOTES.  —  i.  The  ionic  changes  which  occur  during  oxidation  and  re- 
duction are  more  complicated  than  those  of  the  methathetical  reactions 
of  precipitation  and  neutralization ;  in  the  case  of  oxidation  and  reduction 


VOLUMETRIC  ANALYSIS  123 

an  ion  may  change  its  entire  character.    The  electrical  charges  on  the  new 
ion  may  differ  in  sign  as  well  as  in  number  from  those  on  the  original  ion. 
In  equations  expressing  these  ionic  changes  the  algebraic  sum  of  the  charges 
is  always  the  same  on  the  two  sides ;  in  this  case,  for  example,  we  have 
6  Fe+++Cr207— +14  H+=6  Fe++++2  Cr++++7  H20. 

2.  It  is  possible  to  prepare  an  exactly  tenth-normal  solution  of  the 
dichromate  by  dissolving  2.4517  g.  of  the  pure  salt  in  water  and  accurately 
diluting  the  solution  to  500  cc.    The  commercial  salt,  however,  should  not 
be  used  for  this  purpose ;  it  should  be  purified  by  recrystallization  from  hot 
water,  and  then  dried  at  130°. 

3.  The  presence  of  ammonium  sulphate  and  sulphuric  acid  in  the  ferrous 
ammonium  sulphate  seems  to  increase  the  stability  of  the  ferrous  solution. 

4.  The  iron  wire  offered  in  the  market  for  this  purpose  answers  well  as  a 
standard,  and  the  iron  content  of  each  lot  purchased  may  be  ascertained 
by  a  number  of  gravimetric  determinations.     It  may  be  preserved  in  a 
desiccator  over  concentrated  sulphuric  acid,  but  this  must  not  be  allowed 
to  come  in  contact  with  the  wire ;  the  wire  should  always  be  carefully  ex- 
amined for  rust  before  use.     If  necessary,  it  should  be  cleaned  with  fine 
emery  paper. 

5.  The  solution  of  the  wire  in  hot  acid  and  the  short  boiling  insure  the 
removal  of  gaseous  hydrocarbons,  due  to  the  presence  in  the  iron  of  a  small 
amount  of  carbon.    If  not  expelled,  these  might  reduce  some  of  the  dichro- 
mate solution.    Their  complete  expulsion  is  even  more  important  when 
the  wire  is  used  as  a  standard  in  connection  with  potassium  permanganate. 

6.  In  the  determination  of  iron  by  this  method  it  must  be  wholly  present 
in  the  ferrous  condition.    The  common  agents  for  the  reduction  of  ferric 
iron  are  stannous  chloride,  zinc,  sulphurous  acid,  and  hydrogen  sulphide; 
of  these  stannous  chloride  is  the  most  convenient,  but  it  should  be  used  in 
very  slight  excess.    To  this  end,  it  should  be  added  to  the  hot  concentrated 
ferric  solution.    The  removal  of  the  small  excess  of  stannous  chloride  is 
necessary,  and  this  is  readily  accomplished  by  means  of  a  large  excess  of 
mercuric  chloride;  the  chlorides  of  mercury  do  not  react  with  the  iron 
salts  nor  with  the  dichromate  under  the  analytical  conditions.     The  re- 
actions are,  2  FeCl3+SnCl2=2  FeCl2+SnCl4;  and  SnCl2+2  HgCl2=SnCl4 
+Hg2Cl2.    The  mercurous  chloride  is  precipitated. 

It  is  essential  that  stannous  chloride  should  not  be  present  in  great  excess 
and  that  the  solution  should  be  dilute  and  cold  when  the  mercuric  chloride 
is  added;  otherwise  a  secondary  reaction  is  likely  to  take  place  with  the 
reduction  of  mercurous  chloride  to  metallic  mercury  (SnCl2+Hg2Q2=SnCl4 
+  2  Hg) ,  which  would  readily  reduce  the  dichromate  solution.  The  oc- 
currence of  this  secondary  reaction  is  indicated  by  the  darkening  of  the 


124    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

precipitate;  in  such  a  case  the  solution  is  worthless,  and  should  be  dis- 
carded. If  the  ferric  solution  is  hot  and  concentrated  upon  the  addition 
of  the  stannous  chloride,  the  reduction  takes  place  with  the  greatest  ease 
and  only  a  drop  or  two  of  the  stannous  solution  in  excess  need  be  added. 

7.  The  solution  should  be  allowed  to  stand  for  a  minute  or  two  after  the 
addition  of  the  mercuric  chloride,  to  permit  the  complete  precipitation  of 
the  calomel.     It  should  then  be  promptly  titrated  to  avoid  any  reoxidation 
of  the  iron  by  the  air. 

8.  Less  than  30  seconds  are  required  for  the  appearance  of  the  reaction 
with  the  indicator  when  the  ferrous  iron  has  nearly  all  been  oxidized;   if 
the  mixture  is  left  too  long,  the  combined  effect  of  light  and  dust  will  lead 
to  a  partial  reduction  of  the  ferricyanide,  with  the  formation  of  a  blue 
precipitate  of  ferric  ferrocyanide.     Thirty  seconds  is  a  sufficient  interval. 

9.  The  accuracy  of  the  titration  may  be  impaired  by  the  removal  of  too 
much  of  the  solution  for  the  tests ;  for  that  reason  the  tests  should  not  be 
begun  until  most  of  the  iron  has  been  oxidized,  but  at  the  close  of  the  titra- 
tion drops  of  considerable  size  may  properly  be  taken  (see  note  concerning 
this  point  under  the  gravimetric  determination  of  iron).     It  is  best  never 
to  overstep  the  end-point.    The  stirring  rod  should  be  washed  each  time 
in  order  not  to  transfer  any  of  the  indicator  to  the  main  solution.     If  the 
end-point  is  determined  as  prescribed,  the  dichromate  method  is  capable 
of  giving  very  exact  results. 

THE  DETERMINATION  OF  IRON  IN  SIDERITE 

Procedure.  Weigh  out  two  portions  of  about  0.23-0.25  g. 
of  the  finely  powdered  ore  into  300-00.  beakers,  moisten  the 
samples  with  water,  cover  the  beakers,  and  add  to  each  20  cc. 
of  6-normal  hydrochloric  acid  and  about  0.2  g.  of  potassium 
chlorate.  Heat  at  a  temperature  just  below  boiling  until  solvent 
action  has  ceased,  and  to  the  hot  solution  add  stannous  chloride 
solution,  drop  by  drop,  avoiding  an  excess  greater  than  two 
drops.  Add  150  cc.  of  cold  water  and  25  cc.  of  mercuric  chloride 
solution,  allow  to  stand  for  a  minute,  and  proceed  with  the  titra- 
tion as  already  described.  Finally,  calculate  the  percentage 
of  iron  in  the  ore. 

NOTES.  —  i.  Siderite  is  native  ferrous  carbonate;  it  may  contain  some 
organic  matter,  and,  in  order  to  destroy  this,  it  is  directed  to  add  a  little 
potassium  chlorate  to  the  hydrochloric  acid. 


VOLUMETRIC  ANALYSIS  125 

2.  Other  ores  can  of  course  be  analyzed  by  this  method.    Since  most 
of  them  contain  ferric  iron,  and  since  in  the  case  of  ferrous  ores  the  iron  is 
generally  oxidized  during  the  preparation  of  the  solution,  the  quantity  of 
stannous  chloride  required  for  the  reduction  of  the  iron  will  be  much  larger 
than  that  added  to  the  solution  of  iron  wire  in  the  previous  exercise.    In 
no  case,  however,  should  stannous  chloride  solution  be  added  in  greater 
excess  than  two  or  three  drops ;  otherwise  it  is  apt  to  cause  the  separation 
of  metallic  mercury,  and  spoil  the  determination. 

3.  For  another  method  of  dissolving  iron  ores,  see  the  determination  of  iron 
by  means  of  potassium  permanganate,  and  also  Note  i  under  that  method. 

THE  DETERMINATION  OF   CHROMIUM  IN  CHROME 
IRON  ORE 

Procedure.  Weigh  out  two  portions  of  about  0.25  g.  each 
into  iron  crucibles  which  have  been  scoured  inside  until  bright. 
Weigh  out  upon  watch  glasses,  on  the  rough  laboratory  balance, 
two  4-g.  portions  of  dry  sodium  peroxide,  pour  about  three 
quarters  of  each  upon  the  samples  of  ore,  and  mix  the  ore  and 
flux  by  means  of  a  dry  glass  rod.  Remove  any  adhering  par- 
ticles from  the  rod  by  stirring  with  it  the  remaining  peroxide, 
and  then  pour  the  latter  upon  the  surface  of  the  mixture.  (Ow- 
ing to  the  tendency  of  the  peroxide  to  absorb  moisture,  the  first 
portion  should  be  mixed  with  one  sample  before  the  second 
portion  is  weighed  out  from  the  container.)  Place  the  crucible 
upon  a  triangle  and  raise  the  temperature  very  slowly  to  the 
melting  point  of  the  flux,  using  a  low  flame  and  holding  the 
burner  in  the  hand.  Maintain  the  fusion  for  five  minutes, 
stirring  with  a  stout  iron  wire;  do  not  heat  above  moderate 
redness. 

Allow  the  crucible  to  cool  until  it  can  be  held  in  the  hand, 
and  then  cover  it  with  water  in  a  3oo-cc.  beaker,  keeping  the 
beaker  covered  with  a  watch  glass.  When  the  evolution  of 
gas  has  ceased,  rinse  off  and  remove  the  crucible;  heat  the 
solution  to  boiling  for  15  minutes  in  the  covered  beaker,  add 
sufficient  6-normal  sulphuric  acid  (calculated)  to  almost  neutralize 
the  liquid,  and  filter.  To  the  filtrate  and  washings,  which  should 


126     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

be  slightly  alkaline,  add  about  0.2  g.  of  sodium  peroxide,  boil 
for  several  minutes,  acidify  with  6-normal  sulphuric  acid,  and 
add  10  cc.  in  excess. 

Weigh  out  for  each  solution  about  0.5  g.  more  than  enough 
pure  ferrous  ammonium  sulphate  to  reduce  the  chromate,  on 
the  assumption  that  the  ore  was  pure  chromite,  FeO  .  Cr2O3. 
Dissolve  the  salt,  in  a  5oo-cc.  Erlenmeyer  flask,  in  about  50  cc. 
of  freshly  boiled  water  and  20  cc.  of  6-normal  sulphuric  acid, 
and  transfer  to  this  flask  the  chromate  solution,  diluting  the 
whole  to  about  200  cc.  Promptly  titrate  the  excess  of  ferrous 
iron  with  the  standard  dichromate  solution. 

From  the  data  obtained,  calculate  the  percentage  of  chromium 
in  the  ore. 

NOTES.  —  i.  Chrome  iron  ore  consists  essentially  of  ferrous  chromite, 
Fe(Cr02)2.  The  ore  is  decomposed  by  the  flux,  which  oxidizes  the  iron 
oxide  to  sodium  ferrate  and  dissolves  and  oxidizes  the  chromic  oxide  to 
sodium  chromate : 

2  FeO  .  Cr2O3+ioNa2O2=4  Na2CrO4+2  Na2FeO4+4  Na20. 

2.  Fused  sodium  peroxide  attacks  most  materials;  although  it  attacks 
iron  and  nickel,  crucibles  of  these  metals  may  nevertheless  be  used  if  care 
is  taken  to  keep  the  temperature  as  low  as  possible.    The  peroxide  must 
be  dry,  and  no  dust  or  organic  matter  of  any  kind  should  be  present ;  other- 
wise explosions  may  occur. 

3.  When  iron  crucibles  are  used,  the  fusion  should  be  allowed  to  become 
cold   before  it   is   placed   in   water;  otherwise   magnetic   oxide   of   iron, 
FeO  .  Fe203,  is  apt  to  scale  off  from  the  crucible.    This  will  lead  to  no 
error,  however,  if  the  solution  is  only  partially  neutralized  before  filtration. 
Partial  neutralization  is  to  prevent  the  alkali  from  destroying  the  filter 
paper. 

4.  Upon  treatment  with  water  the  chromate  goes  into  solution,  the 
sodium  ferrate  is  decomposed  into  sodium  hydroxide,  ferric  oxide,  and 
oxygen,  and  the  excess  of  sodium  peroxide  is  decomposed  with  the  evolution 
of  oxygen.    The  subsequent  boiling  insures  the  complete  decomposition 
of  the  peroxide,  any  of  which  if  present  would  react  with  the  chromate  upon 
acidification.    The  alkaline  chromate  solution  is  always  slightly  reduced 
upon  filtration  through  a  paper  filter ;  it  is  therefore  directed  to  add  to  the 
filtrate  a  small  quantity  of  sodium  peroxide  to  reoxidize  the  chromium,  and 
to  boil  a  second  time  to  destroy  the  excess. 


VOLUMETRIC  ANALYSIS  127 

5.  The  addition  of  acid  transforms  the  sodium  chromate  into  dichromate, 
which,  of  course,  behaves  like  potassium  dichromate  in  acid  solution.     If 
any  of  the  sodium  peroxide  is  allowed  to  remain  undecomposed  in  the  solu- 
tion, the  chromate  is  at  least  partially  oxidized  to  a  perchromate,  upon 
acidification. 

6.  Instead  of  using  Fe(NH4S04)2 . 6  H2O,  the  ferrous  solution  may  be 
prepared  from  a  suitable  quantity  of  pure  iron  wire ;  or,  of  course,  a  stand- 
ard ferrous  ammonium  sulphate  solution  itself  may  be  added  in  excess. 
Perhaps  an  even  better  method  for  the  determination  of  chromium  consists 
in  the  addition  of  potassium  iodide  in  excess  to  the  acidified  fusion  extract, 
followed  by  the  titration  of  the  iodine  with  sodium  thiosulphate  solution : 

Cr207— +6  I-+i4  H+=  2  Cr++++7  H20+3  I2; 
and  2  Na2S203+l2=  2  NaI+Na2S4O6. 

2.    PERMANGANATE    PROCESSES 

Fundamental  Principles.  In  acid  solution,  potassium  per- 
manganate promptly  and  completely  oxidizes  ferrous  iron  in 
the  cold  to  ferric  iron.  Also,  at  80-90°,  it  reacts  quantitatively 
with  oxalic  acid,  which  it  oxidizes  to  carbonic  acid.  Though  in 
reality  the  reactions  are  not  so  simple,  the  quantitative  relation- 
ships are  accurately  represented  by  the  following  equations : 

10  FeS04+2  KMn04+9  H2S04  =  5  Fe2(S04)3+2  KHS04 

+2MnS04+8H20; 

and  5  H2C2O4-f  2  KMn04+4  H2SO4  =  2  KHSO4+2  MnS04 

+  ioC02+8H20. 

Or,  more  simply  expressed, 

5  Fe+++Mn04-+8  H+  =  5  Fe++++Mn+++4  H2O ; 
and        5  C2O4— +  2  Mn04-+i6  H+  =  2  Mn+++io  CO2+8  H2O. 

In  a  hot  neutral  or  faintly  acid  solution,  in  the  presence  of 
zinc  salts,  potassium  permanganate  oxidizes  manganous  salts 
quantitatively  in  the  sense  of  the  equation, 

3  Mn+++2  Mn04~-f  2  H20  =  4  H++5  Mn02. 

From  these  equations  it  is  readily  seen  that  for  use  in  acid 
solution  the  normal  weight  of  the  salt  is  one  fifth  of  a  mol,  or 
31.61  g.,  while  for  use  in  neutral  solution  the  normal  weight  is 
one  third  of  a  mol,  or  52.68  g. 


128    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

In  addition  to  the  above,  potassium  permanganate  is  capable 
of  oxidizing  stannous,  cuprous,  and  mercurous  salts,  antimonious, 
arsenious,  nitrous,  and  sulphurous  acids,  hydrogen  sulphide, 
ferrocyanides,  and  many  other  substances. 

Furthermore,  as  a  less  desirable  feature,  the  permanganate 
is  capable  under  certain  conditions  of  oxidizing  free  hydrochloric 
acid,  with  the  liberation  of  chlorine;  the  action  is  rapid  in  hot 
or  strongly  acid  solutions,  especially  in  the  presence  of  ferrous 
iron,  but  slow  in  cold  dilute  solutions.  It  is  possible,  however, 
with  suitable  modifications,  to  obtain  very  exact  results  in  the 
presence  of  hydrochloric  acid,  even  in  the  titration  of  iron ;  but, 
other  things  being  equal,  in  acid  solution,  it  is  preferable  to  carry 
out  permanganate  titrations  in  the  absence  of  chlorides. 

Potassium  permanganate  has  an  intense  coloring  power. 
Even  the  tenth-normal  solution  is  so  deeply  colored  that  the 
lower  line  of  the  meniscus  is  not  visible  in  an  ordinary  burette ; 
readings  must  therefore  be  made  from  the  upper  edge.  More- 
over, the  slightest  excess  added  to  an  otherwise  colorless  solu- 
tion is  indicated  with  great  accuracy;  as  its  own  indicator,  it 
renders  the  titration  one  of  the  most  satisfactory  known. 

The  permanganate  solution  should  not  be  placed  in  burettes 
with  rubber  tips ;  it  is  more  or  less  rapidly  reduced  by  most 
organic  substances. 

THE  PREPARATION  AND  STANDARDIZATION  OF  AN  AP- 
PROXIMATELY TENTH-NORMAL  SOLUTION  OF  POTAS- 
SIUM PERMANGANATE 

Procedure.  Dissolve  3.25  g.  of  the  permanganate  crystals 
in  200  cc.  of  warm  water,  dilute  the  solution  to  one  liter,  and  mix 
thoroughly.  The  value  of  this  solution  is  apt  to  change  slowly, 
especially  just  after  it  has  been  prepared.  For  this  reason  the 
solution  should  be  allowed  to  stand  for  several  days,  and  then 
filtered  through  a  layer  of  asbestos  to  remove  the  precipitate 
of  hydrated  manganese  dioxide.  After  thorough  mixing,  it  is 
then  ready  for  standardization.  The  solution  should  be  pre- 


VOLUMETRIC  ANALYSIS  129 

served  in  glass-stoppered  bottles,  and  should  be  protected  from 
heat  and  light.  Thais  prepared  and  preserved,  it  will  retain  its 
oxidizing  value  for  months.  The  solution  is  said  to  be  still  more 
stable  if  it  is  made  very  slightly  alkaline  with  potassium  hy- 
droxide (before  standardization,  of  course). 

Weigh  out  accurately  into  700  cc.  Erlenmeyer  flasks  several 
0.12-0.14  g.  samples  of  pure  sodium  oxalate,  previously  dried 
at  110-120°;  dissolve  each  sample  in  250  cc.  of  hot  water  (80- 
90°),  with  the  addition  of  30  cc.  of  6-normal  sulphuric  acid,  and 
titrate  at  once  with  the  permanganate  solution.  At  first,  the 
permanganate  should  be  added  drop  by  drop,  with  shaking  after 
each  addition  until  the  color  disappears.  After  several  drops 
have  been  added,  the  solution  may  be  run  in  slowly  (10-12  cc. 
per  minute)  with  continuous  shaking.  Toward  the  end  of  the 
titration,  particular  care  must  be  taken  to  allow  the  color  due 
to  each  drop  to  disappear  before  the  addition  of  the  next,  in 
order  to  avoid  passing  the  end-point.  Titrate  to  the  first  per- 
manent pink.  The  temperature  at  the  end  of  the  titration  must 
not  be  below  60°. 

From  the  data  obtained,  calculate  the  normality  factor  of 
the  solution.  Duplicate  values  should  check  within  two  parts 
in  one  thousand. 

NOTES.  —  i.  It  is  not  satisfactory  to  prepare  a  standard  solution  by 
directly  weighing  out  the  calculated  quantity  of  potassium  permanganate, 
even  after  the  latter  has  been  purified  by  recrystallization.  The  best 
practice  is  to  prepare  the  solution  as  described  in  the  procedure,  and  then 
to  standardize  it  by  comparison  with  iron  wire  or  with  sodium  oxalate. 
Ferrous  ammonium  sulphate,  oxalic  acid,  potassium  tetroxalate,  acid  potas- 
sium oxalate,  and  other  substances  have  been  proposed  as  standards,  but 
iron  wire  and  sodium  oxalate  are  readily  obtainable  in  a  sufficiently  pure 
condition,  and  being  non-hygroscopic  and  free  from  water  of  crystallization, 
their  composition  is  less  subject  to  change. 

2.  Upon  treating  a  given  weight  of  pure  sodium  oxalate  with  an  excess 
of  sulphuric  acid,  the  corresponding  weight  of  oxalic  acid  is  set  free;  so 
that  the  use  of  this  salt  as  a  standard  merely  enables  us  easily  to  measure 
out  a  specific  amount  of  oxalic  acid.  The  oxidation  of  the  oxalic  acid  by 
the  permanganate  is  at  first  slow,  and  the  permanganate  should  be  added 


130     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

dropwise,  with  full  time  for  decolorization  between  successive  drops.  After 
a  certain  small  amount  of  manganous  sulphate  has  been  produced  in  the 
solution,  however,  the  speed  of  the  reaction  is  very  greatly  increased  (by 
the  catalytic  action  of  this  substance)  and  the  permanganate  may  be  run 
in  much  faster. 

THE  ^DETERMINATION  OF  IRON  :IN  HEMATITE 

Principles.  One  of  the  most  accurate  methods  for  the  deter- 
mination of  iron  is  based  upon  the  oxidation  of  a  chloride-free 
ferrous  sulphate  solution,  in  the  presence  of  sulphuric  acid,  with 
potassium  permanganate.  Under  these  conditions,  ferrous 
iron  is  oxidized  and  permanganate  is  reduced,  according  to  the 
equation : 

MnO4-+5  Fe+++8  H+  =  Mn+++5  Fe++++4  H20. 

But  if  chlorides  are  present,  some  of  the  permanganate  will 
be  reduced  by  these,  with  the  liberation  (and  partial  escape)  of 
chlorine,  and  the  results  will  be  somewhat  high : 

2  MnO4-+i6  H++IO  Cl-  =  2  Mn+++8  H2O  +  5  C12. 

Upon  the  addition  of  the  permanganate  to  a  cold,  dilute  solu- 
tion of  hydrochloric  acid  alone,  or  to  one  containing  ferric  iron, 
no  chlorine  is  evolved ;  ferrous  iron,  therefore,  seems  to  accel- 
erate this  reaction  by  catalysis. 

Nevertheless,  since  in  dissolving  iron  ores  it  is  nearly  always 
necessary  to  use  strong  hydrochloric  acid,  to  which  it  is  often 
well  to  add  a  little  stannous  chloride,  and  since  stannous  chloride 
is  a  most  convenient  reagent  for  the  reduction  of  ferric  iron  to 
the  ferrous  condition,  it  is  desirable,  if  possible,  to  carry  out  the 
titration  in  the  presence  of  fairly  large  quantities  of  chlorides. 

Now  it  has  been  shown  that  if, 'when  chlorides  are  present,  a 
small  quantity  of  manganous  salt  is  added  to  the  solution,  the 
ferrous  iron  alone  is  oxidized,  and  that  accurate  titrations  can 
be  performed  (Zimmermann) .  But  the  end-point  is  somewhat 
indistinct,  owing  to  the  yellow  tint  of  the  ferric  chloride  pro- 
duced. This  difficulty  can  be  overcome  by  the  addition  of 
phosphoric  and  sulphuric  acids  (Reinhardt),  which  have  recently 


VOLUMETRIC  ANALYSIS  131 

been  shown  to  combine  with  ferric  iron  to  form  colorless  com- 
plexes such  as  H[Fe(S04)2],  H3[Fe(P04)2],  and  H6[Fe(P04)3] 
(Weinland  and  Ensgraber,  Zeitschrift  fur  anorganische  Chemiet 
Vol.  84,  p.  349) ;  the  large  excesses  of  these  acids  repress  the 
dissociation  of  these  complexes  and  insure  a  colorless  solution. 

Procedure.  Weigh  out  three  samples  of  the  finely  ground 
ore,  of  about  0.25  g.  each,  into  100  cc.  beakers.  To  each  sample 
add  15  cc.  of  6-normal  hydrochloric  acid  and  2  cc.  of  stannous 
chloride  solution,  and  gently  heat  the  covered  beakers  for  10-15 
minutes,  until  nothing  other  than  a  small,  white,  sandy  residue 
remains  undissolved.  If  the  hot  solution  is  at  all  yellow,  dis- 
charge this  color  by  adding  stannous  chloride  solution,  one  drop 
at  a  time,  with  stirring ;  avoid  an  excess  of  more  than  two  drops. 
If,  however,  after  the  heating,  the  solution  is  colorless,  stannous 
chloride  is  present  in  unknown  excess,  and  must  be  oxidized 
by  adding  permanganate  solution  (not  to  be  counted,  of  course, 
in  the  volume  required  for  the  titration)  drop  by  drop  with 
stirring,  until  the  yellow  color  due  to  ferric  iron  appears;  dis- 
charge this  color  as  above  directed,  with  stannous  chloride 
solution,  one  drop  in  excess. 

After  cooling,  dilute  the  colorless  solution  with  50  cc.  of  cold 
water,  and  transfer,  with  stirring,  to  a  yoo-cc.  beaker  containing 
10  cc.  of  mercuric  chloride  and  50  cc.  of  water.  (If,  instead  of  a 
white  precipitate  of  calomel,  a  gray  precipitate  of  mercury  is 
formed  at  this  point,  the  solution  must  be  discarded.)  Dilute 
the  mixture  with  cold  water  to  about  500  cc.,  add  8-10  cc.  of 
the  Zimmermann-Reinhardt  solution,1  and  titrate  at  once  with 
the  standard  permanganate  solution.  Add  the  permanganate 
slowly,  with  constant  stirring,  finally  in  single  drops,  until  the 
pink  color  flashes  throughout  the  solution  and  persists  for  15-20 
seconds ;  do  not  pass  the  end-point.  Report  the  percentage  of 
iron  in  the  ore. 

1  Made  by  dissolving  67  g.  of  MnSO4 .  4  H2O  in  500  cc.  of  water,  adding  138  cc. 
of  phosphoric  acid  (sp.  gr.,  1.7)  and  130  cc.  of  sulphuric  acid  (sp.  gr.,  1.84),  and 
diluting  with  water  to  one  liter. 


132     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

NOTES.  —  i.  Many  iron  ores  are  not  completely  decomposed  by  hydro- 
chloric acid,  the  insoluble  residue  containing  more  or  less  iron,  as  silicate, 
titaniferous  iron,  etc.  Unless  iron  is  known  to  be  absent  in  the  insoluble 
residue,  the  finely  ground  sample  should  be  digested  on  the  hot  plate  with 
10  cc.  of  hydrochloric  acid  until  the  residue  is  white,  or  until  there  appears 
to  be  no  further  action;  if  the  ore  contains  carbonaceous  matter,  a  little 
potassium  chlorate  should  be  added.  Finally  evaporate  to  dryness,  ex- 
tract with  5  cc.  of  hydrochloric  acid,  dilute  with  10  cc.  of  water,  allow  to 
settle,  and  decant  the  clear  liquid  through  a  small  filter,  transferring  the 
residue  to  the  filter  and  washing  with  as  little  cold  water  as  possible.  Ignite 
the  filter  and  residue  in  a  small  platinum  crucible,  allow  to  cool,  and  add 
20-30  drops  of  sulphuric  acid  and  twice  as  much  hydrofluoric  acid.  Heat 
carefully,  and,  if  the  residue  is  dissolved,  evaporate  to  white  fumes,  allow 
to  cool,  dissolve  in  water,  and  add  to  the  solution  at  first  obtained.  If, 
however,  this  treatment  fails  to  decompose  the  residue,  drive  off  most  of 
the  sulphuric  acid,  add  0.5-0.6  g.  of  potassium  bisulphate,  and  heat  gradually 
until  the  bisulphate  is  quite  liquid  and  fumes  of  sulphuric  acid  are  given  off 
whenever  the  lid  of  the  crucible  is  raised.  When  the  black  specks  have 
disappeared,  allow  the  crucible  to  cool  and  dissolve  the  salt  in  the  crucible 
with  hot  water  and  a  few  drops  of  hydrochloric  acid. 

In  case  ferric  iron  has  been  dissolved  in  hydrochloric  acid  in  contact  with 
platinum,  the  solution  should  be  oxidized  with  bromine  water  and  the  iron 
precipitated  with  ammonia ;  i.e.  if  it  is  desired  to  use  stannous  chloride  in 
the  reduction.  The  ferric  hydroxide  can  then  be  redissolved  (after  washing 
it  with  hot  water)  in  hydrochloric  acid  and  reduced.  Otherwise  the  iron 
solution  will  contain  a  small  quantity  of  platinum,  4  FeCls+2  HCl+Pt 
=4  FeCl2+H2PtCl6,  which  gives  a  characteristic  ferric-iron  color  with 
stannous  chloride,  and  prevents  the  recognition  of  the  point  at  which  the 
iron  is  reduced. 

2.  Three  samples  should  be  taken,  in  order  that  one  may  be  used  for  a 
rapid  preliminary  titration.    Having  ascertained  in  a  rough  manner  the 
iron  content  oi  the  sample,  the  final  titrations  are  greatly  facilitated. 

3.  Stannous  chloride  is  a  great  help  in  the  solution  of  many  ores  con- 
taining ferric  iron.    Apparently  the  difficultly  soluble  particles  of  hematite 
are  continuously  reduced  at  the  surface  to  ferrous  oxide,  which  is  much 
more  readily  dissolved  by  the  acid. 

4.  The  available  agents  for  the  reduction  of  ferric  iron  are  zinc,  sulphurous 
acid,  and  hydrogen  sulphide ;  stannous  chloride  is  excluded  unless  the  titra- 
tion is  to  be  made  by  the  Zimmermann-Reinhardt  method.    In  that  case 
it  should  be  carefully  added,  in  very  slight  excess,  to  the  hot,  concentrated, 
acid  solution  (cf.  the  standardization  of  dichromate  solution,  Note  6). 


VOLUMETRIC  ANALYSIS  133 

5.  Soluble  salts  of  mercurous  mercury  are  readily  oxidized  by  potassium 
permanganate  in  acid  solution.     Mercurous  chloride,  however,  is  exceed- 
ingly insoluble,  and,  provided  only  a  very  small  quantity  is  suspended  in 
the  solution,  its  action  is  so  slow  that  the  end-point  of  the  titration  can  be 
accurately  fixed.    The  pink  color  which  flashes  throughout  the  solution 
at  the  end  of  the  titration  is,  however,  not  permanent,  and  for  that  reason 
the  time-limit  set  should  be  closely  observed.    For  the  greatest  accuracy, 
the  permanganate  should  of  course  be  standardized,  under  exactly  the  same 
conditions,  against  a  known  quantity  of  metallic  iron.    But  the  error  due 
to  the  use  of  a  solution  standardized  against  sodium  oxalate  is  for  most 
purposes  negligible. 

6.  For  a  rapid  method  for  the  reduction  of  ferric  iron  by  means  of  zinc, 
see  Notes  i  and  2  under  the  Determination  of  Phosphorus  in  Steel.    It 
should  be  noted  that  titanium  is  also  reduced  by  zinc,  but  not  by  the  other 
agents  mentioned ;  with  the  use  of  zinc,  therefore  the  presence  of  titanium 
would  lead  to  high  results. 

THE  DETERMINATION  OF  CALCIUM  IN  LIMESTONE 

Procedure.  Instead  of  igniting  the  precipitate  of  calcium 
oxalate,  obtained  from  the  limestone  by  double  precipitation 
according  to  the  procedure  described  in  Part  II,  and  weighing 
it  as  calcium  oxide,  the  calcium  may  be  determined  volumetrically 
as  follows :  Wash  the  reprecipitated  calcium  oxalate  by  decanta- 
tion,  keeping  it  as  far  as  possible  in  the  precipitation  vessel,  and 
decompose  this  precipitate  by  slowly  pouring  through  the  filter 
at  least  six  5-cc.  portions  of  hot,  3-normal  sulphuric  acid,  wash- 
ing afterwards  with  hot  water,  and  receiving  the  acid  filtrate 
and  washings  in  the  beaker  containing  the  bulk  of  the  precipitate. 
Dilute  this  mixture  to  100  cc.  and  warm  gently,  with  stirring, 
to  completely  decompose  the  calcium  oxalate.  Allow  the  mixture 
to  cool,  transfer  it  quantitatively  to  a  25o-cc.  measuring  flask, 
and  dilute  to  the  mark  with  water,  finally  mixing  the  solution 
by  pouring  it  into  a  clean  dry  beaker  and  back  into  the  flask. 

Measure  out  by  means  of  a  pipette  50.00  cc.  portions  of  this 
solution,  add  to  each  30  cc.  of  6-normal  sulphuric  acid,  dilute  to 
300  cc.,  heat  to  90°,  and  titrate  as  already  described  with  the 
standard  permanganate  solution.  Remembering  that  only  one 


134    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

fifth  of  the  sample  was  used  in  each  titration,  calculate  the  per- 
centage of  CaO  in  the  limestone. 

NOTE.  —  The  reactions  involved  in  the  volumetric  determination  of 
calcium  are:  CaC2O4+H2S04=CaS04+H2C2C>4;  and  5  H2C204+2  KMnO4 
+4H2S04=2KHSO4+2MnS04+ioC02+8H20.  It  is  therefore  plain 
that  the  normal  or  equivalent  weight  of  calcium  oxide  in  this  case  is  one  half 
of  a  mol,  or  that  o.i  N  permanganate  solution  has  a  calcium  oxide  value  of 
0.00280  g.  per  cubic  centimeter. 

THE  DETERMINATION  OF  THE  OXIDIZING  VALUE  OF 
PYROLUSITE 

Procedure.  Weigh  out  two  portions  of  the  very  finely  ground 
mineral,  of  about  0.3  g.  each,  into  500  cc.  Erlenmeyer  flasks. 
Calculate  the  weight  of  ferrous  ammonium  sulphate, 
Fe(NH4S04)2 .  6  H2O,  required  to  react  with  each  sample,  assum- 
ing it  to  be  pure  manganese  dioxide  (2  FeS04+Mn02+2  H2SO4 
=  Fe2(S04)3+MnS04-f-2  H20),  and  weigh  out  accurately  por- 
tions of  the  pure  salt  0.15-0.20  g.  in  excess  of  the  calculated 
amounts,  into  the  corresponding  flasks.  Pour  into  each  flask 
50  cc.  of  water  and  50  cc.  of  6-normal  sulphuric  acid,  cover  the 
flasks,  and  heat  to  boiling  until  the  action  is  complete.  Finally, 
dilute  to  about  300  cc.,  and  promptly  titrate  the  excess  of  fer- 
rous iron  with  the  standard  permanganate  solution.  From  the 
data  obtained,  calculate  the  percentage  of  Mn02  in  the  sample. 

NOTES.  —  i.  The  mineral  should  be  so  finely  ground  that  no  grit  what- 
ever can  be  detected  when  a  little  of  the  powder  is  placed  between  the 
teeth ;  upon  this  the  success  of  the  analysis  largely  depends.  If  properly 
ground,  solution  will  be  complete  in  10-15  minutes. 

2.  A  moderate  excess  of  ferrous  iron  is  necessary  to  promote  rapid  solu- 
tion, and,  also  to  facilitate  solution,  the  mixture  should  not  be  diluted  be- 
fore the  solvent  action  has  ceased. 

3.  A  solution  of  iron  wire  in  sulphuric  acid  may  be  substituted  for  the 
ferrous  ammonium  sulphate,  but  in  that  case  there  is  more  danger  of 
the  partial  oxidation  of  the  iron  by  the  air.    For  example,  if  iron  wire  is 
used,  it  should  be  dissolved  in  sulphuric  acid  out  of  contact  with  air,  and 
the  air  should  not  have  access  to  the  solution  during  cooling.    This  is  best 
accomplished  by  means  of  a  Contat-Gockel  valve,  which  consists  of  a  glass 


VOLUMETRIC  ANALYSIS  135 

bulb  with  an  inner  siphon,  as  shown  in  the  figure.    In  the  bulb  is  placed  a 

cold  saturated  solution  of  sodium  bicarbonate,  through  which  the  hydrogen 

(and  steam)  evolved  in  the  flask  bubbles.    After  all 

the  iron  has  been  dissolved,  the  liquid  is  boiled  for  a 

few  minutes  longer,  and  the  flame  is  removed.    As  the 

flask  cools  off,  small  portions  of  the  bicarbonate  are 

at  intervals  sucked  into  the  flask  and  decomposed  by 

the  acid  with  the  evolution  of  carbon  dioxide,  whereby 

the  entrance  of  more  bicarbonate  solution  is  prevented. 

For  other  methods  of  performing  this  analysis,  see 
Part  IV,  Problems  23,  73  and  74.  According  to  O.  L. 
Barnebey  (J.  Ind.  Eng.  Chem.,  Vol.  9,  p.  961  (1917)), 
the  use  of  oxalic  acid  in  place  of  the  ferrous  salt  yields 
less  reliable  results. 

4.  With  the  substitution  of  very  dilute  nitric  acid 
for  sulphuric  acid  in  the  above  procedure,  the  method 
may  be  used  to  determine  the  oxidizing  power  of  red 
lead,  or  minium,  Pb304,  and  of  lead  peroxide,  Pb02. 
Of  these  substances,  samples  of  i.o  and  0.8  g.,  respectively,  should  be  taken 
when  3o-cc.  burettes  are  used.    It  is  better,  however,  to  make  use  of  an 
iodometric  method. 


THE  DETERMINATION  OF  PHOSPHORUS  IN  STEEL 

Principle.  The  molybdic  anhydride  contained  in  ammonium 
phosphomolybdate,  (NH4)3P04 .  12  MoO3,  may  be  reduced  by 
zinc  in  the  presence  of  sulphuric  acid,  from  Mo03  to  Mo203; 
but  molybdenum  in  the  latter  condition  is  not  stable  in  the 
presence  of  air.  If,  however,  the  acidified  molybdate  solution 
is  passed  through  a  Jones  reductor  (see  below)  directly  into  a 
solution  of  ferric  sulphate,  the  sensitive  molybdic  compound  is 
oxidized  by  the  ferric  salt  with  the  formation  of  an  equivalent 
amount  of  ferrous  sulphate,  less  sensitive  to  the  atmospheric 
action.  The  molybdenum  solution  is  green  as  it  leaves  the 
reductor,  but  upon  mixing  with  the  ferric  salt  the  green  color 
disappears;  if  phosphoric  acid  is  added,  the  color  due  to  the 
presence  of  ferric  iron  is  destroyed.  The  decolorized  solution 
is  titrated  while  still  hot  with  tenth-normal  permanganate 


136     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

solution,  of  which  the  quantity  necessary  corresponds  to  the 
equation, 

5  Mo2O3+6  KMn04  =  3  K20+6  MnO+io  Mo03. 
From  this  it  may  be  seen  that  5  P  =c=  30  Mo203  036  KMn04  ~  90  0, 
or  Poi8  0;  one  cubic  centimeter  of  o.i  N  permanganate  solu- 
tion represents,  therefore,  0.0862  mg.  of  phosphorus. 

Procedure.  Weigh  out  two  samples  of  steel  drillings,  each 
sufficient  to  contain  1.7-2.0  mg.  of  phosphorus,  into  250-0:. 
Erlenmeyer  flasks.  Add  to  each  a  mixture  of  25  cc.  of  nitric 
acid  (sp.  gr.,  1.42)  and  75  cc.  of  water.  Suspend  in  the  neck 
of  each  flask  a  small  funnel  and  heat  until,  after  complete  solu- 
tion, the  oxides  of  nitrogen  have  been  expelled.  Dissolve  0.3- 
0.4  g.  of  KMn04  crystals  in  10  cc.  of  hot  water,  add  one  half  of 
this  solution  to  the  contents  of  each  flask,  and  boil  until  the 
permanganate  color  has  disappeared.  Remove  the  flame,  add 
sulphurous  acid  or  ammonium  bisulphite  solution,  a  few  drops 
in  excess,  to  dissolve  the  precipitated  oxides  of  manganese,  boil 
out  the  excess  of  sulphur  dioxide,  and  filter  the  solution;  re- 
ceiving the  filtrate  in  a  similar  flask.  Add  ammonia  to  the 
solution  with  stirring  until  a  permanent  precipitate  just  begins 
to  form,  and  then  add  nitric  acid  drop  by  drop  to  clear  up  the 
solution.  Finally,  at  a  temperature  of  40°,  add  40  cc.  of  molyb- 
date  solution,  close  the  flask  with  a  rubber  stopper,  and  shake 
vigorously  for  five  minutes;  allow  the  precipitate  to  settle. 
(At  this  point,  prepare  the  Jones  reductor  for  use,  as  described 
in  Note  2.) 

Now  filter  the  solution,  keeping  the  precipitate  as  far  as  pos- 
sible in  the  flask,  and  wash  by  decantation  with  a  solution  of 
ammonium  sulphate  acidified  with  sulphuric  acid l  until  the 
washings  give  no  test  for  molybdenum  with  ammonium  sulphide 
and  hydrochloric  acid.  Dissolve  the  precipitate  by  pouring 
through  the  filter  a  mixture  of  5  cc.  of  6-normal  ammonia  and 
20  cc.  of  water,  and  collecting  the  filtrate  and  washings  in  the 

1  Made  by  mixing  15  cc.  of  ammonia  (sp.  gr.,  0.90)  and  25  cc.  of  sulphuric  acid 
(sp.  gr.,  1.84)  with  one  liter  of  water. 


VOLUMETRIC  ANALYSIS 


137 


precipitation  flask.  Acidify  the  solution,  which  should  have  a 
volume  of  about  60  cc.,  with  10  cc.  of  sulphuric  acid  (sp.  gr., 
1.84)  and  promptly  pass  the  acidified  solution,  before  it  has  a 
chance  to  cool  off,  through  the  reductor  into  the  receiver  (collect- 
ing the  liquid  beneath  the  surface  of  100  cc.  of  a  solution  con- 
taining 25  g.  of  ferric  alum  and  40  cc.  of  sirupy  phosphoric  acid, 
sp.  gr.,  1.7,  per  liter),  preceded  by  100  cc.  of  hot  water  and 
followed  by  200  cc.  of  hot  dilute  sulphuric  acid  (i  :  40)  and  100  cc. 
of  hot  water.  See  that  no  air  enters  the  reductor  during  this 
entire  operation.  Titrate  the  reduced  solution  at  once  with 
tenth-normal  permanganate,  and  calculate  the  percentage  of 
phosphorus  in  the  steel  on  the  assumption  that  the  yellow  pre- 
cipitate contains  phosphorus  and  molybdenum  in  the  proportion 
indicated  by  the  formula  (NH^sPOi  .  12  MoOa. 


190mm. 


NOTES.  —  i.  The  Jones  reductor,  which  also  is  useful  in  the  reduction 
of  ferric  iron  for  titration,  is  essentially  a  column  of  amalgamated  zinc, 
through  which  the  solution  is 
passed  for  reduction.  It  is  as- 
sembled as  shown  in  the  accom- 
panying figure.  The  tube  A  has 
an  inside  diameter  of  about  18  mm. 
and  (for  this  reduction)  is  400  mm. 
long  ;  the  small  extension  tube  has 
an  inside  diameter  of  6  mm.  and  a 
length  of  300  mm.  below  the  stop- 
cock. At  the  base  of  the  tube  A 
are  placed  some  glass  beads  ;  these 
are  covered  by  a  plug  of  glass  wool 
several  millimeters  thick,  and  upon 
this  is  placed  a  layer  of  asbestos, 
such  as  is  used  for  Gooch  filters, 
not  exceeding  i  mm.  in  thickness. 
The  tube  is  then  filled  with  the 
amalgamated  zinc  to  within  50  mm. 
of  the  top,  and  this  is  covered  with 
a  plug  of  glass  wool.  The  reductor  is  connected  as  shown  with  the  suction 
bottle  F,  and  the  bottle  D  is  a  safety  vessel  to  prevent  contamination  of  the 
solution  from  the  suction  apparatus. 


138     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  amalgamated  zinc  is  prepared  by  dissolving  5-6  g.  of  mercuric  chlo- 
ride in  250  cc.  of  water,  with  the  addition  of  5-10  cc.  of  dilute  hydrochloric 
acid,  adding  to  this  solution  500  g.  of  (20-30  mesh)  granulated  zinc,  in  a  large 
flask,  and  shaking  thoroughly  for  two  minutes ;  the  solution  is  then  poured 
off  and  the  zinc  thoroughly  washed  with  water. 

2.  To  prepare  the  reductor  for  use,  connect  the  safety  bottle  with  the 
vacuum  pump,  fill  the  reductor  while  the  stopcock  is  nearly  closed  with 
warm,  dilute  sulphuric  acid  (25  cc.  of  the  concentrated  acid  in  one  liter), 
and  then  open  the  stopcock  so  that  the  acid  runs  through  slowly.     Continue 
to  pour  in  acid  until  200-300  cc.  have  passed  through,  then  close  the  cock 
while  some  liquid  is  still  left  in  the  funnel.     (During  the  whole  operation, 
see  that  no  air  enters  the  reductor;  if  air  enters,  hydrogen  peroxide  will 
be  formed  from  oxygen  and  nascent  hydrogen,  and  the  results  will  be  worth- 
less.)   Now  remove  the  filtrate,  and  again  pass  through  200  cc.  of  the 
warm  acid,  followed  by  100  cc.  of  warm  water;  test  this  liquid  (300  cc.) 
with  the  standard  permanganate  solution,  in  order  to  determine  the  volume 
of  permanganate  required  to  color  the  acid  alone.    This  amount  must  be 
subtracted  from  the  volume  required  in  the  subsequent  titration. 

3.  Upon  dissolving  the  steel  in  nitric  acid  of  the  strength  indicated,  the 
phosphorus  is  oxidized,  and  none  of  it  is  lost  by  evolution  as  phosphine. 
The  permanganate  is  subsequently  added  in  order  to  insure  the  complete 
oxidation  of  carbonaceous  matter  and  of  the  phosphorus  to  phosphoric  acid. 

4.  The  higher  oxides  of  manganese,  as  Mn02,  are  not  soluble  in  nitric 
acid.    Upon  the  addition  of  a  reducing  agent,  however,  such  as  hydrogen 
peroxide  or  sulphurous  acid,  their  solution  is  effected : 

MnO2+H2S03=MnO+H2SO4~MnS04+H2O. 

5.  In  connection  with  the  precipitation  of  phosphoric  acid  as  ammonium 
phosphomolybdate,  the  student  should  consult  the  notes  under  the  Deter- 
mination of  Phosphoric  Anhydride. 

6.  Since  the  molybdenum  in  the  precipitate  prepared  from  one  gram  of 
a  steel  containing  0.15%  of  phosphorus  would  require  by  this  method 
17.44  cc.  of  o.i -normal  permanganate  solution,  it  is  readily  seen  that  the 
process  is  a  rapid  one  for  arriving  at  very  accurate  results.    This  is  es- 
pecially true  if  the  permanganate  has  been  standardized  under  the  same 
conditions  against  a  steel  of  accurately  known  phosphorus  content;  in 
such  a  case,  it  would  be  unnecessary  to  correct  for  the  small  amount  of  iron 
extracted  from  the  (impure)  amalgamated  zinc,  since  this  would  be  the 
same  in  both  standardization  and  analysis. 

7.  It  scarcely  needs  to  be  pointed  out  that  the  method  is  not  suitable  for 
determining  phosphorus  or  phosphoric  acid  in  substances  containing  them 


VOLUMETRIC  ANALYSIS  139 

in  large  amount.  This  would  require  for  titration  relatively  enormous 
quantities  of  permanganate  solution,  and,  what  is  still  worse,  it  would  be 
practically  impossible  to  completely  reduce  the  molybdenum.  If  small 
aliquot  portions  were  taken  for  reduction  and  titration,  any  error  of  meas- 
urement would  be  multiplied  by  a  very  large  factor  in  the  calculation  of 
the  result. 

8.  The  following  method  is  suitable  for  the  volumetric  determination  of 
phosphorus  or  phosphoric  acid  when  these  are  present  in  larger  amounts. 
The  phosphorus  or  phosphoric  acid  is  converted  into  ammonium  phospho- 
molybdate ;  this,  after  washing  with  KNO3  solution,  is  dissolved  in  an  excess 
of  standard  sodium  hydroxide  solution ;  and  the  resulting  solution  is  titrated 
with  standard  nitric  acid,  with  phenolphthalein  as  an  indicator.  Needless 
to  say,  the  sodium  hydroxide  should  be  standardized  under  identical  condi- 
tions against  a  sample  of  accurately  known  phosphorus  content. 

THE  DETERMINATION  OF   MANGANESE  IN  AN  ORE 

Fundamental  Principles.  When  potassium  permanganate  is 
added  to  a  hot,  neutral  or  very  faintly  acid  solution  of  manganese 
sulphate  a  reaction  takes  place  (the  Guyard  reaction)  in  which 
the  manganous  oxide  of  the  sulphate  is  oxidized  at  the  expense 
of  the  anhydride  of  the  permanganate,  with  the  precipitation 
of  hydrated  intermediate  oxides  in  varying  proportions.  These 
are  manganous  acid,  MnO(OH)2,  and  hydrated  salts  of  man- 
ganous acid.  The  essential  changes  in  the  state  of  oxidation 
may  be  represented  as  follows : 

Mn207+3  MnO  =  5  Mn02 ; 
Mn207+8  MnO  =  5(Mn02 .  MnO) ; 
Mn207+i3  MnO  =  5(Mn02 .  2  MnO). 

It  is  clear,  then,  that  in  this  form  the  Guyard  reaction  cannot 
furnish  the  basis  for  a  satisfactory  volumetric  method. 

It  has  been  found,  however,  that  under  suitable  conditions, 
in  the  presence  of  zinc  ion,  a  hydrated  manganite  of  zinc  is  pre- 
cipitated, which,  while  variable  in  composition,  contains  all  the 
manganese  in  the  quadrivalent  condition.  Thus  regulated,  the 
reaction  furnishes  a  valuable  means  for  the  determination  of 
manganese  (Volhard's  Method).  Although  the  composition  of 


140     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  precipitate  varies,  the  course  of  the  reaction  is  typically 

represented  by  the  following  equation : 

4  KMn04+5  ZnS04+6  MnS04+i4  H20  =  2  K2S04+9  H2S04 

+5Zn(oMn/°H), 
or,  more  simply, 

2  MnO4~+3  Mn+++2  H20=4  H++5  Mn02. 

Procedure.  Weigh  out  into  a  5oo-cc.  Erlenmeyer  flask  a 
sufficient  quantity  of  the  very  finely  ground  ore  to  contain  about 
0.20  g.  of  manganese;  add  3  g.  of  potassium  chlorate  and  20  cc. 
of  i2-normal  hydrochloric  acid,  and  boil  until  the  ore  is  com- 
pletely decomposed  and  the  chlorine  expelled.  Dilute  with 
water  to  about  50  cc. ;  transfer  the  cold  solution  quantitatively 
to  a  loo-cc.  measuring  flask;  dilute  to  the  mark  with  water; 
and  mix  thoroughly  by  pouring  the  contents  of  the  flask  into 
a  clean,  dry  beaker,  and  back  into  the  flask. 

Now,  from  a  burette  or  pipette,  measure  into  500-0:.  Erlen- 
meyer flasks  four  20.00  cc.  portions  of  this  solution,  and  treat 
each  as  follows :  Dilute  with  water  to  100  cc.,  heat,  and  to  the 
acid  solution  add  with  shaking  an  aqueous  suspension  of  zinc 
oxide,1  in  small  portions,  until  the  iron  is  completely  precipitated 
as  ferric  hydroxide ;  this  point  may  be  recognized  by  the  sudden 
coagulation  of  the  precipitate,  upon  shaking,  and  the  decoloriza- 
tion  of  the  brownish  colored  solution.  The  precipitate  should 
not  be  light  yellow,  but  should  have  the  characteristic  brownish 
red  color  of  ferric  hydroxide,  and  the  least  possible  excess  of 
zinc  oxide  should  be  used.  (Should  the  ore  contain  a  quantity 

1  Dissolve  100  g.  of  crystallized  zinc  sulphate  in  300  cc.  of  hot  water,  and  with 
stirring  cautiously  add  to  the  clear  solution  a  few  drops  of  a  solution  made  by 
dissolving  25-27  g.  of  pure  sodium  hydroxide  in  150  cc.  of  water,  until  the  zinc 
solution  remains  distinctly  turbid ;  then  add  a  little  bromine  water,  heat,  and  filter. 
To  the  nitrate  add  the  bulk  of  the  sodium  hydroxide  solution,  and  stir.  Rinse  the 
mixture  into  a  one-liter  bottle,  and  fill  the  latter  with  water.  The  mixture  should 
be  well  shaken  when  used.  (This  suspension  should  not  react  alkaline  with  phenol- 
phthalein,  and  a  10  cc.  portion  of  the  mixture,  when  cleared  up  with  sulphuric  acid, 
diluted  to  100  cc.,  and  treated  with  one  drop  of  o.i  N  KMnO4,  should  be  perma- 
nently colored  pink.) 


VOLUMETRIC  ANALYSIS  141 

of  iron  insufficiently  in  excess  of  that  required  by  any  phosphoric 
and  arsenic  acids  present,  then  5  cc.  of  a  solution  containing 
20  g.  of  ferric  chloride  per  liter  should  be  added  before  the  pre- 
cipitation with  zinc  oxide.)  If  too  much  zinc  oxide  is  added, 
the  solution  will  be  milky ;  in  that  case  very  dilute  hydrochloric 
acid  should  be  added  drop  by  drop  to  the  hot  solution  until  the 
supernatant  liquid  just  becomes  clear. 

Finally  dilute  the  solutions  to  300  cc.  and,  at  80°,  treat  them 
successively  as  follows :  Run  into  the  first  solution  the  standard 
permanganate  in  5  cc.  portions,  until  after  continued  shaking 
the  liquid  retains  a  permanent  pink  tinge,  —  say  after  the  addi- 
tion of  the  fifth  portion  (i.e.  25  cc.) ;  into  the  second  solution 
run  5  cc.  less  permanganate  than  the  volume  previously  used 
(e.g.  20  cc.),  shake  until  the  pink  color  disappears,  and  then  finish 
the  titration  by  the  further  addition  of  permanganate  in  portions 
of  i  cc.  until  the  pink  color  persists  after  protracted  shaking, 
say  after  23.0  cc.  in  all  have  been  added;  to  the  third  solution 
add  at  once  i.o  cc.  less  permanganate  than  the  total  volume  used 
in  the  second  case  (e.g.  22.0  cc.),  and  continue  the  titration  with 
the  addition  of  0.20  cc.  portions,  until  the  hot  solution  matches 
in  color  a  solution  prepared  by  the  addition  of  o.io  cc.  of  the 
permanganate  to  300  cc.  of  water.  With  the  fourth  solution, 
repeat  this  titration.  If,  for  example,  22.60  cc.  of  the  perman- 
ganate have  been  used  in  each  of  the  last  two  titrations,  then 
this  quantity  minus  the  o.io  cc.  of  the  solution  used  for  com- 
parison should  be  taken  as  the  volume  actually  required.  Report 
the  percentage  of  manganese  in  the  ore. 

NOTES.  —  i.  In  case  the  treatment  with  hydrochloric  acid  and  potas- 
sium chlorate  should  be  insufficient  to  thoroughly  decompose  the  ore  (in- 
dicated by  the  presence  of  a  dark-colored  residue),  the  residue  should  be 
filtered  off,  washed,  dried,  and  ignited  hi  a  platinum  crucible.  It  should 
then  be  fused  with  sodium  carbonate,  the  melt  dissolved  in  hydrochloric 
acid,  and  the  solution  evaporated  to  dryness  in  a  porcelain  dish,  to  dehy- 
drate the  silica.  The  final  residue  should  be  moistened  with  hydrochloric 
acid,  taken  up  in  water,  and  filtered  into  the  Erlenmeyer  flask  containing 
the  acid  filtrate  from  the  original  residue.  The  resulting  solution,  which 


142     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

contains  all  the  manganese,  is  then  evaporated  to  a  small  volume,  trans- 
ferred to  the  measuring  flask,  and  treated  as  described  in  the  procedure. 

2.  Upon  the  addition  of  zinc  oxide  to  the  acid  solution  of  the  ore,  the 
zinc  oxide  first  neutralizes  the  acid  with  the  formation  of  zinc  chloride,  and 
then  precipitates  the  iron  with  the  further  formation  of  zinc  chloride,  accord- 
ing  to   the  reaction:   2  FeCl3+3  ZnO+3  H2O  =  2  Fe(OH)3+3  ZnCl2.    In 
this  way  sufficient  zinc  ion  is  introduced  into  the  solution  to  insure  the 
conversion  of  the  manganese  into  the  hydrated  manganite  of  zinc. 

3.  Although  it  is  often  recommended  to  convert  the  chlorides  in  the 
solution  into  sulphates  before  the  addition  of  zinc  oxide,  this  treatment  is 
not  necessary.    The  titration  of  manganese  in  a  dilute  neutral  solution 
with  potassium  permanganate  is  a  very  different  thing  from  that  of  ferrous 
iron  in  a  dilute  acid  solution  containing  chlorides.    In  the  latter  case,  the 
ferrous  iron  catalyzes  the  reaction,  2  KMn04+i6  HC1=2  KC1+2  MnCl2 
+  5  Cl2-|-8  H20,  some  of  the  chlorine  escapes,  and  there  is  a  tendency  to 
high  results.    In  the  former  case,  however,  nothing  is  present  in  the  solution 
to  catalyze  the  reaction  between  the  permanganate  and  the  small  quantity 
of  hydrochloric  acid  which  is  formed ;  and,  although  the  solution  is  hot,  its 
acid  concentration  is  so  low  that  there  is  no  danger  from  this  source.     Start- 
ing with  0.2000  g.  of  an  ore  containing  20%  of  manganese,  for  example, 
the  total  quantity  of  acid  formed  in  the  titration  (e.g  4  KMn04+  5  ZnCl2 
+6  MnCl2+i4  H20=4  KCl+i8  HCl+s  ZnO  .  Mn203(OH)2)  weighs  about 

TT/"M 

^r— — Xo.04,  or  somewhat  less  than  o.i  g. ;  and  this  quantity  in  a  volume 

of  over  300  cc.  would  give  an  acid  strength  of  less  than  o.oi-normal. 

4.  The  titration  should  be  performed  at  80-85°,  and  especial  care  should 
be  taken  not  to  heat  the  solution  too  hot  during  the  titration. 

5.  For  the  greatest  accuracy,  in  spite  of  all  that  has  been  said  above, 
the  permanganate  solution  should  be  standardized  against  a  known  quantity 
of  manganese,  weighed  as  MnS04,  under  conditions  similar  to  those  to  be 
used  in  the  analysis. 


3.    IODOMETRIC   PROCESSES 

Fundamental  Considerations.  Analytical  processes  which 
depend  upon  the  volumetric  measurement  of  specific  amounts 
of  iodine  are  known  as  iodometric  methods.  In  these  processes, 
either  iodine  is  used  in  standard  solution  to  bring  about  a  definite 
reaction,  or  the  iodine  liberated  in  a  reaction  is  determined  by 
titration  with  some  suitable  standardized  reagent. 


VOLUMETRIC  ANALYSIS  143 

The  titration  of  iodine  against  sodium  thiosulphate,  with 
starch  as  an  indicator,  is  one  of  the  most  accurate  of  volumetric 
processes.  The  process  may  be  used  in  neutral  or  slightly  acid 
solutions  to  determine  free  iodine,  and  this  in  turn  may  serve 
as  a  measure  of  any  substance  capable  of  liberating  iodine  from 
hydriodic  acid.  For  example,  the  quantity  of  potassium  iodate 
in  a  sample  of  the  salt  may  be  determined  on  the  basis  of  the 
reactions  : 

KI03+6  KI+6  HC1  =  6  KCl+KI+3  I2+3  H20; 
and  I2+2  Na2S203  =  2  NaI+Na2S406. 

It  should  be  noted  that  chlorine  and  bromine  oxidize  sodium 
thiosulphate  partially  to  sulphate,  while,  under  analytical  con- 
ditions, iodine  oxidizes  it  wholly  to  sodium  tetrathionate. 

Iodine  acts  as  an  oxidizing  agent  either  through  the  direct 
withdrawal  of  a  positive  constituent,  as  shown  in  the  equations  : 
2Na2S203+I2  =  2NaI+Na2S406,  and  H2S+I2  =  2  HI+S,  or 
through  the  decomposition  of  water  in  the  presence  of  a  reducing 
agent,  as  in  the  equations:  H2S03+I2+H20$H2S04+2  HI, 
and  H3As03+I2+H20  =  H3As04+2  HI  (H2O+I2:£HI+HOI; 
and  H3As03+HOI^lH3As04+HI).  It  will  be  seen  from  these 
equations  that  a  one-tenth  normal  iodine  solution  contains  one 
tenth  of  one  gram-atom,  or  12.692  g.  of  iodine  per  liter. 

The  solubility  of  iodine  in  water  is  too  small  for  the  prepara- 
tion of  even  a  one-tenth  normal  solution.  In  the  presence  of 
sufficient  potassium  iodide,  however,  the  iodine  dissolves  much 
more  readily,  owing  to  the  formation  of  an  unstable  but  soluble 
polyiodide  of  the  formula  KI3  : 

KI  .  I2,  or  I-+I2^(I  .  I,)-. 


In  the  presence  of  reducing  agents  iodine  is  removed  from  this 
equilibrium  mixture,  the  reaction  runs  to  completion  from  right 
to  left,  and  the  solution  can  be  used  as  though  it  were  a  simple 
solution  of  iodine.  The  potassium  iodide  used  in  the  prepara- 
tion of  the  solution  should  weigh  about  1.5  times  as  much  as 


144     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  iodine.  Moreover,  the  presence  of  potassium  iodide  in  the 
solution  renders  it  possible  to  employ  commercial  iodine  (which 
is  apt  to  contain  chlorine  as  an  impurity)  in  the  preparation  of 
the  standard  solutions;  the  chlorine  is  removed  according  to 
the  equation,  IC1+KI  =  KC1+I2. 

In  performing  iodometric  titrations  in  the  presence  of  sul- 
phuric acid,  particular  attention  should  be  given  to  the  main- 
tenance of  suitable  analytical  conditions.  If,  for  example,  it  is 
desired  to  determine  copper  by  titrating  the  iodine  liberated  in 
the  reaction,  2  CuS04+2  H2S04+4  KI$Cu2I2+4  KHS04+I2, 
it  is  not  sufficient  to  simply  add  potassium  iodide  and  sul- 
phuric acid  in  (unknown)  excess.  It  must  be  remembered 
that  such  a  mixture  will  contain  both  sulphuric  and  hydriodic 
acid,  and  that  if  the  concentration  of  either  is  too  great,  or  if 
the  solution  is  allowed  to  become  at  all  warm,  the  determination 
is  very  apt  to  be  spoiled:  H2S04+2  HI  =  H2S03+H2OH-I2; 
H2S04+6HI  =  S+4H20-{-3l2;  or,  in  extreme  cases,  H2S04 
+8HI  =  H2S+4H20-|-4 12.  Other  things  being  equal,  when 
acid  solutions  are  required,  it  is  better  to  use  acetic  acid  or  dilute 
hydrochloric  acid. 

In  direct  titrations  with  iodine,  e.g.  in  the  presence  of  sodium 
bicarbonate,  it  is  best  to  work  in  the  absence  of  ammonium 
salts.  Such  solutions  are  very  faintly  alkaline,  especially  if  at 
all  warm;  and  in  the  presence  of  ammonium  salts  ammonia 
is  apt  to  be  liberated,  which  is  not  entirely  without  influence 
upon  the  titration. 

Iodine  solutions  act  upon  rubber,  so  that  burettes  with  glass 
stopcocks  should  be  used. 

Determination  of  the  End-point.  A  single  drop  of  one-tenth 
normal  iodine  solution  imparts  a  distinct  tint  to  200  cc.  of  water, 
and  in  many  titrations  with  this  solution  no  other  indicator  is 
required.  If,  however,  the  solution  to  be  titrated  contains  colored 
substances,  or  if  the  greatest  possible  accuracy  is  demanded,  a 
solution  of  starch  should  be  used  as  an  indicator.  Under  the 
proper  conditions,  the  presence  of  one  part  of  free  iodine  in 


VOLUMETRIC  ANALYSIS  145 

several  millions  of  solution  can  be  recognized  with  this  indi- 
cator, but  the  sensitiveness  of  the  reaction  and  the  color  pro- 
duced are  affected  by  a  number  of  factors.  The  test  is  decidedly 
more  sentitive  when  the  concentration  of  iodide  ion  (and  of  hydro- 
gen ion l)  is  not  too  low,  and  when  the  quantity  of  starch  present 
is  sufficient  to  give  a  deep  blue  color. 

Under  less  favorable  conditions,  the  starch  may  give  a  greenish 
or  a  reddish  color ;  or  it  may  be  very  unreliable,  as  in  solutions 
containing  an  abnormally  low  iodide  ion  concentration.  How- 
ever, since  the  standard  iodine  solution  always  contains  potas- 
sium iodide,  and  since  an  iodide  is  always  one  product  of  the 
titration,  there  is  ordinarily  not  much  danger  from  this  source. 
Attention  should  be  directed  mainly  toward  the  observance  of 
uniform  conditions  in  all  related  titrations :  the  volumes  of  the 
solutions  titrated  should  be  approximately  equal,  the  starch 
solution  should  be  properly  prepared,  and  the  same  quantity  of 
it  should  be  added  for  each  titration.  Finally,  all  titrations 
should  be  made  in  the  cold ;  the  iodo-starch  blue  is  discharged 
by  heat. 

Preparation  of  the  Starch  Solution.  Rub  i  g.  of  potato  starch 
with  5  cc.  of  cold  water  to  a  smooth  paste,  and  slowly  add  this 
to  200  cc.  of  boiling  water.  Continue  the  boiling  for  about  i 
min.  until  an  almost  clear  solution  is  obtained,  set  this  aside  to 
settle,  and  finally  decant  the  supernatant  liquid  through  a  filter. 
Use  5  cc.  of  the  clear  filtrate  for  each  titration. 

A  "  soluble  starch  "  which  is  in  the  market  is  more  convenient, 
since  with  it  filtration  is  unnecessary.  A  solution  made  by  add- 
ing 200  cc.  of  boiling  water  to  i.o  g.  of  this  starch,  previously 
mixed  with  a  little  cold  water,  serves  the  purpose  well.  Use 
5  cc.  of  this  solution  for  each  titration. 

In  either  case,  the  starch  solution  should  be  freshly  prepared. 
If  a  great  many  titrations  are  to  be  made,  however,  it  is  ad- 
visable to  prepare  a  liter  of  the  starch  solution;  a  number  of 

1  The  iodo-starch  blue  is  discharged  by  caustic  alkalies  and  somewhat  less  readily 
by  sodium  or  potassium  carbonate,  but  not  by  the  bicarbonates. 
L 


146    INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

small  (50-100  cc.)  bottles  should  be  filled  with  this  solution, 
heated  for  two  hours  in  a  water  bath,  and,  while  still  in  the  bath, 
they  should  be  closed  with  paraffined  soft  cork  stoppers.  Thus 
sterilized  and  protected  from  the  air,  the  solution  will  retain 
its  sensitiveness  almost  indefinitely.  After  a  bottle  has  been 
opened,  mould  nearly  always  begins  to  form  within  a  few  days ; 
hence  the  use  of  small  bottles. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROX- 
IMATELY ONE-TENTH  NORMAL  SOLUTIONS  OF  IODINE 
AND  SODIUM  THIOSULPHATE 

Procedure.  Weigh  out  on  the  rough  laboratory  balance 
6.3-6.4  g.  of  commercial  iodine,  add  it  to  a  solution  of  9  g.  of 
potassium  iodide  in  25  cc.  of  water,  in  an  Erlenmeyer  flask,  and 
agitate  the  mixture  until  the  iodine  is  completely  dissolved. 
Dilute  the  solution  to  500  cc.,  in  a  measuring  flask,  and  mix  it 
thoroughly. 

Heat  600-700  cc.  of  distilled  water  in  a  large  flask  and  boil  for 
about  5  minutes.  Stopper  the  flask  loosely,  and  allow  the  water  to 
cool.  Weigh  out  1 2.5  g.  of  sodium  thiosulphate,  Na2S203 .  5  H2O, 
introduce  it  into  a  5oo-cc.  measuring  flask,  and  dissolve  it  in 
about  200  cc.  of  the  cold,  freshly  boiled  water.  Finally  dilute  to 
the  mark  with  more  of  the  same  water,  and  mix  thoroughly. 

After  these  solutions  have  come  to  the  room  temperature, 
fill  a  burette  with  each  (the  iodine  in  a  glass-stoppered  burette), 
observing  the  usual  precautions  to  prevent  dilution.  Run  out 
25  cc.  of  the  thiosulphate  solution  into  a  beaker,  dilute  with 
150  cc.  of  water,  add  5  cc.  of  starch  solution,  and  titrate  with 
the  iodine  to  the  appearance  of  the  blue  of  the  iodo-starch.  If 
the  end-point  is  overstepped,  titrate  back  with  the  thiosulphate 
solution.  (All  waste  solutions  containing  iodine  and  potassium 
iodide  should  be  poured  into  the  vessel  provided  for  iodine  residues.) 
Repeat  until  the  ratio  of  the  two  solutions  is  accurately  estab- 
lished, taking  into  account  all  necessary  corrections  for  burettes 
and  for  temperature  changes. 


VOLUMETRIC  ANALYSIS  147 

Standardization  of  the  Iodine  Solution.  Weigh  out  into 
500-cc.  beakers  two  o.i2-o.i3-g.  portions  of  pure  arsenious  oxide, 
and  in  each  case  dissolve  the  arsenious  oxide  in  10  cc.  of  6-normal 
sodium  hydroxide  solution.  Dilute  the  solution  to  100  cc., 
add  2  drops  of  methyl  orange,  and  then  cautiously  add  6-normal 
hydrochloric  acid  until  the  solution  contains  2  or  3  drops  in 
excess.  Cover  the  beakers,  add  to  each  a  solution  of  5  g. 
of  pure  sodium  bicarbonate  in  75  cc.  of  cold  water,  then  add 
5  cc.  of  starch  solution,  and  titrate  with  the  iodine  to  the  ap- 
pearance of  the  blue  color.  Do  not  pass  the  end-point.  From 
the  corrected  volume  of  the  iodine  solution  used,  calculate  the 
normality  factor  of  the  solution.  Duplicate  values  should  agree 
within  two  parts  in  one  thousand.  Also,  from  the  ratio  pre- 
viously found,  calculate  the  normality  factor  of  the  thiosulphate 
solution. 

NOTES.  —  i.  Iodine  solutions  are  acted  upon  by  sunlight  with  the 
formation  of  hydriodic  acid,  and  a  high  room-temperature  tends  to  volatilize 
the  iodine.  They  require  frequent  standardization  against  pure  arsenious 
acid,  anhydrous  sodium  thiosulphate,  or  standard  thiosulphate  solution. 

2.  Sodium  thiosulphate,  Na2S203  .  5  H2O,  may  be  obtained  pure  by 
recrystallization.    It  is  then  possible  to  prepare  a  standard  solution  by 
dissolving  the  calculated  weight  of  the  salt  in  pure  cold  water,  and  dilution 
to  the  required  volume.     Such  solutions  are  quite  stable  and  may  be  kept 
for  months  without  appreciable  change  in  concentration,  provided  they  are 
not  allowed  to  absorb  carbon  dioxide  ;  but  they  should  be  protected  from 
heat  and  light,  both  of  which  are  likely  to  promote  decomposition. 

3.  Carbonic  acid  causes  a  slow  decomposition  of  the  thiosulphate  solu- 
tion, with  the  formation  of  free  sulphur  and  sulphurous  acid;  and,  since 
sulphurous  acid  acts  in  the  same  way,  the  decomposition  once  started  be- 
comes progressive  : 

xro  Q  n  -LTT  rn  ->  I  Na2C03+ 
Na2S203+H2C03;±  |  H2Sa03 

Na2S03+ 


The  reducing  value  of  the  solution  increases  gradually  as  the  decomposition 
progresses  ;  i.e.  the  solution  apparently  becomes  stronger.  When  it  is  con- 
sidered that  in  this  decomposition  each  molecule  of  thiosulphate  yields  one 


148     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

molecule  of  sulphite^the  greater  reducing  value  is  readily  understood ;  for 
iNa&Qi+Iii-NftiSjOrHNal,  while  iNa2SO8+I2+H2O=Na2S04-|-2HL 

4.  Solutions  of  the  thiosulphate  may  be  standardized  against  pure 
iodine,  or,  more  conveniently  (with  the  help  of  pure  potassium  iodide), 
against  potassium  bromate,  potassium  iodate,  or  potassium  dichromate. 
These  three  salts  are  readily  obtainable  in  a  pure  condition. 

5.  Arsenious  oxide  dissolves  most  readily  in  caustic  alkalies,  and  for  this 
reason  the  sodium  hydroxide  is  used.    The  presence  of  sodium  hydroxide 
is  not  admissible,  however,  during  the  titration,  since  it  reacts  readily  with 
iodine.    It  is  therefore  removed  by  the  addition  of  a  slight  excess  of  acid, 
and  sodium  bicarbonate  is  then  added  in  large  excess.    The  purpose  of  the 
bicarbonate,  which  under  the  analytical  conditions    is  without   action 
upon  the  iodine,  is  to  neutralize  the  acid  formed  in  the  reversible  reaction, 
As203-f-2  I2+2  H20^rAs205-|-4  HI,  and  thus  cause  it  to  run  to  completion 
from  left  to  right.    The  reaction  may  then  be  written : 

Na2HAs03+I2+2  NaHC03=Na2HAs04+2  NaI+2  C02+H20. 

6.  Since  the  addition  of  iodine  in  excess  to  the  weakly  basic  bicarbonate 
solution  is  likely  to  lead  to  a  slight  degree  of  action,  it  is  best  in  this  titra- 
tion not  to  overstep  the  end-point. 

7.  Iodine  is  a  rather  expensive  chemical  and  it  is  well  worth  while  to 
recover  it  from  the  united  residues  of  a  large  class. 

THE  DETERMINATION  OF  ANTIMONY  IN  STIBNITE 

The  sample  for  analysis  should  be  an  antimony  ore,  practically 
free  from  arsenic  and  iron,  and  with  hydrochloric  acid  it  should 
leave  only  a  siliceous  residue. 

Procedure.  Weigh  out  two  0.20  g.  portions  of  the  finely 
ground  mineral  into  dry  i5o-cc.  beakers.  Cover  the  beakers, 
add  5  cc.  of  i2-normal  hydrochloric  acid,  and  allow  the  acid 
to  act  in  the  cold  for  10  minutes;  then  heat  gently  on  the 
steam  bath,  for  about  15  minutes,  until  the  residue  is  white. 
Add  2  g.  of  powdered  tartaric  acid  and  gently  warm  the  mixture 
for  10  minutes  longer.  Do  not  allow  the  liquid  to  evaporate 
sufficiently  to  expose  any  part  of  the  bottom  of  the  beaker. 
Dilute  the  solution  cautiously  with  5-cc.  portions  of  water; 
if  a  red  coloration  appears,  stop  the  dilution,  warm  until  the 
solution  is  colorless,  and  again  dilute.  Continue  the  dilution 


VOLUMETRIC  ANALYSIS  149 

until  a  volume  of  100  cc.  is  reached,  and  boil  for  a  minute. 
Neutralize  the  clear,  cold  solution  with  sodium  hydroxide 
(methyl  orange),  and  then  acidify  it  with  dilute  hydrochloric 
acid,  a  drop  or  two  in  excess. 

Dissolve  10  g.  of  sodium  bicarbonate  in  400  cc.  of  water,  place 
200  cc.  of  this  solution  in  each  of  two  yoo-cc.  beakers,  and  trans- 
fer to  these  the  cold  solutions  of  the  ore,  avoiding  loss  by  effer- 
vescence. Add  5  cc.  of  starch  solution,  and  titrate  each  mixture 
with  the  standard  iodine  solution,  to  the  appearance  of  the  blue 
color.  Do  not  overstep  the  end-point. 

From  the  corrected  data,  calculate  the  percentage  of  antimony 
in  the  stibnite. 

NOTES.  —  i.  Stibnite  is  essentially  native  antimony  sulphide,  Sb2S8, 
and  upon  treatment  with  hydrochloric  acid  hydrogen  sulphide  is  liberated ; 
but  this  is  partially  absorbed  by  the  acid.  The  gas  should  be  wholly  ex- 
pelled during  the  heating  on  the  steam  bath ;  if  it  is  not  completely  driven 
out,  antimony  sulphide  will  begin  to  separate  at  some  point  in  the  dilution. 
In  that  case,  however,  if  the  dilution  is  at  once  stopped  and  the  solution 
heated,  the  hot  acid  will  redissolve  the  sulphide,  and  the  hydrogen  sulphide 
may  then  be  expelled.  The  final  boiling  is  to  insure  the  absence  of  hydro- 
gen sulphide,  which  itself  reacts  with  iodine. 

2.  Antimony  trichloride  hi  the  presence  of  strong  hydrochloric  acid  is  some- 
what volatile,  and  for  this  reason  the  solution  should  not  be  boiled  before 
dilution.    If  the  solution  is  gently  heated  as  described,  no  error  need  be 
feared  from  this  source. 

3.  If,  for  any  reason,  a  white  precipitate  of  oxy chloride  separates  during 
the  dilution  or  neutralization,  it  is  best  to  reject  the  solution  and  start 
anew.    Antimony  chloride  is  readily  hydrolyzed  upon  dilution,  with  the 
precipitation  of  basic  compounds  such  as  SbOCl;  but  the  addition  of  tar- 
taric  acid  leads  to  the  formation  of  stable  antimonyl  tartrates,  which  are 
soluble.    In  this  way  the  antimony  is  kept  in  solution. 

4.  The  reaction  between  the  iodine  and  the  antimonyl  tartrate  is  not 
so  simple,  but  for  purposes  of  calculation  it  is  accurately  expressed  by  the 
equation,  Sb203+2  I2+2  H20^±Sb205+4  HI.    The  purpose  of  the  bicar- 
bonate is  here  also  to  neutralize  the  hydriodic  acid  formed,  and  thereby 
drive  the  oxidation  to  completion. 

5.  The  sodium  hydroxide  is  added  merely  to  neutralize  most  of  the  acid, 
and  to  make  it  easy  to  provide  for  the  presence  of  a  known  amount  of 


150     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

sodium  bicarbonate.    The  solution  should  be  slightly,  but  distinctly,  acid 
when  the  bicarbonate  is  added. 

6.  If  the  ore  to  be  analyzed  contains  more  than  traces  of  iron,  it  is  dis- 
solved in  hydrochloric  acid,  the  antimony  is  precipitated  with  hydrogen 
sulphide,  and  the  washed  precipitate  is  redissolved  in  hydrochloric  acid 
and  determined  as  above.  In  case  arsenic  also  is  present,  a  somewhat 
more  complicated  separation  of  the  antimony  is  necessary. 

THE  DETERMINATION  OF  LEAD  IN  AN  ORE 

Procedure.  Weigh  out  two  samples  of  the  finely  ground  ore 
sufficient  to  contain  about  0.20  g.  of  lead  (0.28-0.29  g-  °f  a  7°% 
ore),  and  treat  each  as  follows :  Moisten  the  sample  with  water, 
add  15  cc.  of  i2-normal  hydrochloric  acid,  and  evaporate  on 
the  steam  bath  to  about  5  cc.  Add  3  cc.  of  strong  nitric  acid, 
evaporate  nearly  to  dryness,  then  add  20  cc.  of  6-normal  hydro- 
chloric acid  and  again  heat  to  bring  all  the  lead  chloride  into 
solution.  Add  20  cc.  of  6-normal  sulphuric  acid  and  evaporate 
to*  white  fumes.  Allow  to  cool,  add  50  cc.  of  water,  boil,  and 
then  add  15  cc.  of  alcohol ;  stir,  allow  to  settle,  and  filter.  Wash 
the  lead  sulphate  and  gangue  six  times  with  lo-cc.  portions  of 
o.5-normal  sulphuric  acid  (15  cc.  of  6-normal  acid  in  165  cc.  of 
water),  transfer  the  residue  to  a  small  beaker  by  means  of  a  jet 
of  water,  and  heat  it  gently  for  a  few  minutes  with  20  cc.  of 
ammonium  acetate  solution ; l  filter  the  liquid  through  the  orig- 
inal filter  and  wash  the  latter  with  small  portions  of  the  hot 
ammonium  acetate  solution.  Dilute  the  extract  to  150  cc., 
heat  to  boiling  and  add  from  a  pipette  10  cc.  of  potassium  dichro- 
mate  solution.2  Boil  the  mixture  gently  for  10  minutes,  filter 
off  the  precipitate  of  lead  chromate  and  wash  the  filter  and 
precipitate  about  ten  times  with  lo-cc.  portions  of  dilute  am- 
monium acetate  solution  (25  cc.  of  the  extraction  solution  diluted 
to  250  cc.),  until  the  excess  of  potassium  chromate  is  completely 
removed. 

1  Made  by  neutralizing  30%  acetic  acid  with  6-normal  ammonia,  and  then  add- 
ing a  slight  excess  of  ammonia. 

2  A  solution  containing  75  g.  of  K2Cr2O7  per  liter. 


VOLUMETRIC  ANALYSIS  151 

Now  place  a  clean  5oo-cc.  Erlenmeyer  flask  under  the  funnel, 
and  with  a  jet  of  cold,  acid,  sodium  chloride  solution  1  stir  up 
and  dissolve  the  precipitate ;  continue  washing  with  the  same 
liquid  until  every  trace  of  color  is  removed  from  the  filter.  In 
any  case,  use  at  least  50  cc.  of  the  liquid.  Finally  dilute  to 
150  cc.,  add  i  g.  of  potassium  iodide,  mix,  and  titrate  at  once 
with  a  solution  of  sodium  thiosulphate  (which  has  been  stand- 
ardized in  the  same  way  against  test  lead,  see  Note  5)  until 
the  brown  color  becomes  faint ;  then  add  5  cc.  of  starch  solution, 
and  continue  the  titration  cautiously  until  the  solution  becomes 
pale  green  (CrCl3)  with  no  tinge  of  blue.  The  end-point  is  very 
sharp,  but  without  great  care  it  may  easily  be  passed.  It  is 
best  to  have  a  white  surface  under  the  flask. 

Report  the  percentage  of  lead  in  the  ore. 

NOTES.  —  i.  The  ore  is  first  heated  with  strong  hydrochloric  acid  in 
order  to  expel  most  of  the  sulphur.  Nitro-hydrochloric  acid  is  then  used 
to  decompose  any  refractory  sulphides.  Upon  evaporating  the  chloride 
solution  to  white  fumes  with  sulphuric  acid,  the  volatile  acids  in  which  lead 
sulphate  is  slightly  soluble  are  completely  expelled,  and  upon  dilution  with 
water,  especially  if  alcohol  is  added,  the  lead  is  all  left  in  the  residue  as  lead 
sulphate. 

2.  Lead  sulphate  is  readily  dissolved  by  ammonium  acetate  solution, 
owing  to  the  exceptional  behavior  of  lead  acetate  with  respect  to  ionization 
(see  Part  I),  leaving  the  siliceous  gangue,  BaS04,  etc.,  as  a  residue. 

3.  While  lead  is  not  precipitated  from  solutions  containing  a  large  ex- 
cess of  acetate  ion  by  sulphates,  the  addition  of  a  soluble  chromate  causes 
the  precipitation  of  lead  chromate.    This  behavior  is  due  to  the  fact 
that  such  solutions  contain  Pb++-ion  at  an  extremely  low  concentration 
(owing  to  the  presence  of  the  lead  mainly  in  the  form  of  intermediate  or  com- 
plex ions,  as  (Pb  .  C2H302)+,  [Pb(C2H302)3]~,  etc.),  and  also  to  the  fact  that 
lead  sulphate  is  very  much  more  soluble  than  lead  chromate ;  the  lead-ion 
concentration  is  still  great  enough  in  such  solutions  to  cause  the  solubility 
product  of  lead  chromate  to  be  exceeded  upon  the  addition  of  potassium 
chromate  in  excess. 

4.  The  solubility  of  lead  chromate  in  the  acid  chloride  solution  is  due 
on  the  one  hand  to  the  lowered  concentration  of  the  chromate  ion,  owing  to 

1  Mix  10  cc.  of  i2-normal  hydrochloric  acid  with  15  cc.  of  water,  and  add  this 
mixture  to  100  cc.  of  a  saturated  solution  of  sodium  chloride. 


152     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  formation  of  non-ionized  H2Cr04,  HCr04~,  etc.,  and  on  the  other  hand 
to  the  great  tendency  of  lead  ion  to  form  soluble  complexes  with  chloride 
solutions.  (Cf.  the  solubility  of  silver  chloride  in  chloride  solutions.) 

5.  The  reaction  of  the  acid  solution  with  potassium  iodide  is  most  simply 
represented  by  the  equation, 

Cr2Of-+6  I-+I4  H+=  2Ci++N-7  HzO+3  I2; 


from  this  it  may  be  seen  that  one  atom  of  lead  (as  PbCr04)  leads 
to  the  liberation  of  3  atoms  of  iodine.  But  the  results  vary  slightly 
with  the  conditions,  and  for  that  reason  the  thiosulphate  solution  must 
be  standardized  under  identical  conditions  against  a  known  amount  of 
lead.  In  this  case,  0.20  g.  of  test  lead  should  be  dissolved  in  5  cc.  of 
6-normal  nitric  acid,  the  solution  evaporated  to  white  fumes  with  20  cc. 
of  6-normal  sulphuric  acid,  and  the  subsequent  operations  carried  out  as 
described  in  the  procedure. 

THE  DETERMINATION  OF   COPPER  IN  AN  ORE 

Principle.  This  method  is  based  upon  the  reaction  which 
takes  place  upon  the  addition  of  potassium  iodide  to  a  slightly 
acid  copper  salt  solution;  cuprous  iodide  is  precipitated  as  a 
cream-colored  powder,  and  iodine  is  set  free  : 

2  CuS04+4  KI  =  2  K2SO4+Cu2l2+l2. 

The  iodine  is  promptly  titrated  with  a  standard  thiosulphate 
solution. 

Standardization  of  the  Thiosulphate  Solution.  Weigh  ac- 
curately two  portions  of  pure  bright  copper  wire  or  foil,  of  0.15- 
0.16  g.  each,  and,  in  25o-cc.  Erlenmeyer  flasks,  dissolve  these  in 
5-cc.  portions  of  6-normal  nitric  acid.  Dilute  each  solution  to 
15  cc.  and  boil  to  expel  the  red  fumes;  then  dilute  to  25  cc.  and 
add  ammonia  (sp.  gr.,  0.90)  in  slight  excess.  Again  boil  until 
the  ammonia  odor  is  faint,  add  80%  acetic  acid,  2-3  cc.  in  excess, 
and  boil  for  a  moment  longer,  agitating  the  flask  in  a  holder  to 
prevent  bumping.  Cool  to  room  temperature,  dilute  to  40  cc., 
add  a  solution  of  3  g.  of  potassium  iodide  in  10  cc.  of  water,  and 
titrate  at  once  with  the  approximately  tenth-normal  thiosul- 
phate solution  to  a  faint  brown  tinge  ;  add  5  cc.  of  starch  solu- 


VOLUMETRIC  ANALYSIS  153 

tion,  and  continue  the  titration  until  the  last  faint  lilac  tint  is 
removed  by  a  single  drop.  Do  not  overstep  the  end-point.  From 
the  data  obtained,  calculate  the  value  of  the  solution  per  cubic 
centimeter  in  terms  of  copper.  $«•> 

Analytical  Procedure.  Weigh  out  into  3oo-cc.  beakers  samples 
of  the  ore  sufficient  to  furnish  about  0.15  g.  of  copper,  and  treat 
each  as  follows:  Add  10  cc.  of  hydrochloric  acid  (sp.  gr.,  1.19) 
and  5  cc.  of  nitric  acid  (sp.  gr.,  1.42)  and  heat  in  the  covered 
beaker  on  the  hot  plate  until  decomposition  is  complete,  adding 
more  of  the  acids  if  necessary,  and  enough  water  at  the  end  to 
hold  all  soluble  salts  in  solution.  Then  add  15  cc.  of  6-normal 
sulphuric  acid,  and  continue  the  heating  until  abundant  white 
fumes  begin  to  come  ofl.  Cool,  add  50-60  cc.  of  water,  boil  for 
a  moment,  and  allow  to  stand,  hot,  until  any  anhydrous  ferric 
sulphate  has  dissolved.  Finally,  filter  off  from  any  lead  sulphate, 
gangue,  and  sulphur,  receiving  the  filtrate  and  washings  in  a 
3Oo-cc.  beaker.  Now  add  a  solution  of  5  g.  of  sodium  thio- 
sulphate  in  25  cc.  of  water,  boil  to  coagulate  the  precipitate, 
and  filter,  transferring  the  precipitate  quantitatively  to  the 
filter  by  means  of  hot  water.  Dry  the  precipitate  on  the  filter. 

Place  the  precipitate,  together  with  the  filter,  in  a  porcelain 
crucible,  ignite  gently  until  the  filter  is  consumed,  and  allow  to 
cool.  Transfer  the  bulk  of  the  precipitate  to  a  2  5o-cc.  Erlenmeyer 
flask,  and  set  aside.  To  dissolve  the  last  portions  of  the  pre- 
cipitate from  the  crucible,  add  3  cc.  of  concentrated  nitric  acid 
and  2  cc.  of  water,  and  warm  gently  on  the  hot  plate,  finally 
pouring  the  acid  solution  into  the  flask  containing  the  bulk  of 
the  precipitate,  and  washing  out  the  crucible  with  a  few  small 
portions  of  6-normal  nitric  acid.  Heat  the  mixture  in  the  flask 
until  the  decomposition  is  complete,  dilute  to  25  cc.,  boil,  add 
ammonia  in  slight  excess,  and  heat  until  the  odor  is  fault.  Add 
80%  acetic  acid,  2-3  cc.  in  excess,  and  boil  for  a  moment,  vigor- 
ously agitating  the  flask  to  prevent  bumping.  Cool  to  room 
temperature,  dilute  to  40  cc.,  add  3  g.  of  potassium  iodide  dis- 
solved in  10  cc.  of  water,  and  titrate  at  once  with  the  thio- 


154     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

sulphate  solution,  as  previously  described.     Report  the  per- 
centage of  copper  in  the  ore. 

NOTES.  —  i.  Since  iron  and  other  elements  likely  to  be  present  inter- 
fere with  the  process,  the  copper  must  be  separated  from  these.  Lead  is 
first  removed  by  means  of  sulphuric  acid,  after  which  the  copper  is  pre- 
cipitated from  the  hot,  acid  solution  by  means  of  sodium  thiosulphate ; 
this  gives  a  flocculent  precipitate  of  cuprous  sulphide  mixed  with  sulphur, 
which  filters  readily  and  can  be  washed  with  hot  water  without  fear  of 
oxidation.  Arsenic  and  antimony,  if  present,  are  also  precipitated,  but 
under  the  treatment  prescribed  the  usual  quantities  of  these  elements  are 
without  influence.  They  are  mostly  volatilized  during  the  ignition.  If 
antimony  is  present  in  appreciable  quantity,  it  is  perhaps  better  to  filter 
the  solution  before  the  addition  of  the  ammonia. 

2.  In  order  to  obtain  the  best  results  it  is  necessary  to  standardize  the 
thiosulphate  solution  against  pure  metallic  copper.    When  this  is  done  the 
method  is  very  accurate ;  otherwise  the  results  are  not  so  good.    For  ex- 
ample, a  thiosulphate  solution  which  (titrated  against  a  freshly  stand- 
ardized iodine  solution)  had  a  calculated  copper  value  of  0.00608  g.  per 
cubic  centimeter,  was  found  upon  standardization  against  pure  copper  to 
have  a  value  of  0.006  n  g.  per  cubic  centimeter. 

3.  Since  nitrous  fumes  liberate  iodine  from  potassium  iodide,  they  must 
be  completely  expelled  by  boiling  before  the  addition  of  the  salt.    The 
expulsion  of  the  last  traces  of  these  fumes  is  insured  by  boiling  the  solu- 
tion after  it  has  been  acidified  with  acetic  acid. 

4.  The  return  of  the  blue  tinge  in  the  liquid  after  long  standing  is  of  no 
significance,  but  a  quick  return  which  is  not  prevented  from  recurring  by 
the  addition  of  a  single  drop  of  the  thiosulphate  solution  is  usually  an 
evidence  of  faulty  work. 

5.  In  such  a  case,  or  if  the  end-point  has  accidentally  been  passed,  the 
same  sample  may  be  prepared  anew  for  titration :  Add  10  cc.  of  concen- 
trated nitric  acid,  and  heat  very  cautiously,  with  great  care  not  to  allow 
the  mixture  to  foam  over.    After  most  of  the  iodine  has  been  expelled, 
manipulate  the  flask  (in  a  holder)  over  a  free  flame  and  boil  the  solution 
down  rapidly  to  a  volume  of  5-10  cc.    Dilute  to  25  cc.  with  water,  boil, 
add  ammonia  in  slight  excess,  and  finish  as  described  in  the  procedure. 

6.  In  the  electrolytic  determination  of  copper  in  ores  containing  arsenic 
and  other  interfering  substances,  a  satisfactory  copper  solution  is  most 
readily  prepared  by  dissolving  the  ignited  thiosulphate  precipitate  in  a 
suitable  quantity  of  strong  nitric  acid,  with  subsequent  dilution  to  the 
required  volume. 


VOLUMETRIC  ANALYSIS  155 

C.  PRECIPITATION  METHODS 

General  Discussion.  The  completion  of  neutralization  re- 
actions depends  upon  the  very  slight  degree  of  ionization  of  one 
of  the  products,  water.  The  completion  of  reactions  of  oxida- 
tion and  reduction  most  often  depends  upon  the  relative  poten- 
tials of  oxidizing  and  reducing  agents  under  specific  experi- 
mental conditions.  Certain  other  reversible  reactions  which 
serve  as  the  basis  of  volumetric  processes  run  to  completion  in 
consequence  of  the  formation  of  very  slightly  soluble  precipi- 
tates. In  most  cases  an  indicator  is  used,  but  in  some  the  cessa- 
tion of  precipitation  with  the  further  addition  of  the  standard 
solution  indicates  the  completion  of  the  reaction. 

An  example  of  the  latter  kind  is  found  in  Gay-Lussac's  method 
for  silver,  which  dates  from  1832,  and  which  is  still  widely  used 
in  determining  the  fineness  of  silver  bullion.  When  silver 
chloride  first  separates  it  is  finely  divided,  and  a  very  minute 
quantity  can  easily  be  recognized;  if  the  solution  is  shaken 
vigorously  the  precipitate  coagulates  and  settles,  leaving  a 
supernatant  liquid  which  is  perfectly  bright  and  clear.  Hence, 
if  silver  nitrate  is  titrated  with  a  solution  of  sodium  chloride 
and  the  mixture  well  shaken  in  a  stoppered  bottle  after  each 
addition,  the  point  at  which  the  addition  of  a  further  quantity 
of  the  standard  solution  fails  to  produce  a  precipitate  can  readily 
be  determined.  Near  the  end-point  it  is  customary  to  use  a 
standard  solution  ten  times  as  dilute  as  that  used  at  the  start. 

In  the  case  of  this  reaction,  this  method  of  determining  the 
end-point  admits  of  a  very  high  degree  of  accuracy,  and  it  is 
the  method  in  use  at  the  government  mints.  Since,  however, 
it  is  rather  tedious  and  demands  considerable  skill  and  experi- 
ence, a  slightly  less  accurate  but  much  more  convenient  method 
is  generally  employed. 

Silver  thiocyanate  is  even  less  soluble  than  silver  chloride, 
and  it  is  therefore  possible  to  titrate  silver  very  accurately  with 
a  standard  solution  of  potassium  or  ammonium  thiocyanate.  If  a 


156     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

solution  of  a  ferric  salt,  acidified  to  suppress  hydrolysis,  is  present 
as  an  indicator,  the  first  drop  of  thiocyanate  solution  in  excess 
will  impart  to  the  mixture  a  pink  tint.  This  method  (Volhard's) 
is  also  suitable  for  the  determination  of  the  halogens  (except 
fluorine)  and  of  certain  other  ions  which  give  silver  compounds 
insoluble  in  dilute  nitric  acid.  A  measured  volume  of  standard 
silver  nitrate  solution  is  added  in  excess,  and  the  excess  is  then 
determined  by  means  of  the  standard  thiocyanate  solution. 

THE  PREPARATION  AND  STANDARDIZATION  OF  APPROX- 
IMATELY ONE-TENTH  NORMAL  SOLUTIONS  OF  SILVER 
NITRATE  AND  POTASSIUM  THIOCYANATE 

Procedure.  Dissolve  5.0  g.  of  potassium  thiocyanate  (or  the 
equivalent  quantity  of  the  ammonium  salt)  in  water  and  dilute 
the  solution  to  500  cc.  Also  dissolve  8.5  g.  of  silver  nitrate  in 
water  and  dilute  the  solution  to  500  cc.  Further,  mix  10  cc. 
of  6-normal  nitric  acid  with  40  cc.  of  water,  heat  the  solution  to 
boiling,  and  dissolve  in  the  hot  liquid  5  g.  of  pure  ferric  alum. 
Allow  the  solution  to  cool,  and  keep  it  for  use  as  an  indicator. 

Now  fill  the  burettes  with  the  respective  solutions,  placing 
the  silver  nitrate  solution  in  a  glass-stoppered  burette.  (Ob- 
serve the  usual  precautions  to  prevent  dilution,  and  place  all 
solutions  and  precipitates  containing  silver  in  the  receptacle  for 
silver  residues.)  Run  out  20  cc.  of  the  silver  nitrate  into  a  beaker, 
dilute  to  150  cc.,  add  10  cc.  of  6-normal  nitric  acid  which  has  been 
recently  boiled,  and  5  cc.  of  the  indicator  solution.  Run  in 
the  thiocyanate  solution  until,  after  vigorous  stirring,  a  faint 
pink  tinge  can  be  detected  in  the  solution.  If  the  end-point  is 
overstepped,  titrate  back  with  the  silver  nitrate  solution.  From 
the  corrected  volumes  used,  calculate  the  ratio  of  the  thiocyanate 
to  the  silver  nitrate  solution.  Repeat  until  the  results  do  not 
differ  by  more  than  two  parts  in  one  thousand. 

Finally,  standardize  the  silver  nitrate  solution,  as  follows: 
Weigh  out  portions  of  pure  sodium  chloride,  of  0.12-0.14  g. 
each,  dissolve  these  in  75-cc.  portions  of  water,  heat  to  boiling, 


VOLUMETRIC  ANALYSIS  157 

and  with  stirring  run  into  each  from  a  burette  25.00  cc.  of  the 
silver  nitrate  solution.  Add  10  cc.  of  freshly  boiled  6-normal 
nitric  acid,  stir,  and  filter,  washing  the  precipitate  by  decantation 
with  several  small  portions  of  hot  distilled  water,  and  pouring 
these  slowly  over  the  filter;  the  united  filtrate  and  washings 
should  have  a  volume  of  about  150  cc.  To  this  solution  add 
5  cc.  of  the  indicator,  and  titrate  the  excess  of  silver  with  the 
thiocyanate  solution,  as  already  described.  From  the  data 
obtained,  calculate  the  normality  factor  of  the  silver  nitrate 
solution;  and  from  the  mean  of  the  duplicate  values,  which 
should  agree  within  two  parts  in  a  thousand,  calculate  the  nor- 
mality factor  of  the  thiocyanate  solution. 

NOTES.  —  i.  The  reactions  between  the  thiocyanate  and  the  indicator 
are  essentially  as  follows  : 

Fe++++6  CNS-3*Fe(CNS).+3 


It  will  be  recalled  that  in  testing  for  ferric  iron  with  potassium  thiocyanate, 
it  is  necessary  to  add  a  large  excess  of  the  latter  in  order  to  detect  the  smallest 
possible  quantity  of  iron.  In  the  same  way,  when  using  ferric  iron  as  an 
indicator  for  thiocyanate,  it  is  necessary  to  provide  a  high  concentration 
of  the  former  in  order  to  detect  the  slightest  possible  excess  of  the  thio- 
cyanate in  the  solution.  The  reactions  which  give  rise  to  the  colored  sub- 
stances are  reversible,  but  in  the  presence  of  a  large  excess  of  one  of  the 
colorless  constituents  the  dissociation  of  the  colored  substances  is  prevented 
by  mass  action. 

2.  Nitric  acid  is  added  to  the  solution  to  be  titrated  in  order  to  prevent 
the  hydrolysis  of  the  ferric  salt,  which  would  impart  a  brownish  red  color 
to  the  mixture.    It  is  boiled  to  free  it  from  nitrous  fumes,  though  this  is  of 
less  importance  here  than  in  testing  for  iron  in  the  presence  of  nitric  acid  ; 
nitrou?  fames  color  the  thiocyanate  pink. 

3.  Sodium  chloride  may  easily  be  obtained  pure  by  filtering  a  concen- 
trated solution  of  the  commercial  salt,  saturating  it  with  hydrogen  chloride 
gas,  and  filtering  off  the  precipitate.    The  latter  is  washed  with  strong 
hydrochloric  acid  and  dried  at  150°,  or  higher. 

4.  Standard  solutions  of  silver  nitrate  can  of  course  be  prepared  by  the 
solution  of  the  calculated  amount  of  pure  metallic  silver  in  nitric  acid,  and 
dilution  to  the  required  volume;  or  by  means  of  the  calculated  weight  of 
pure  silver  nitrate. 


158     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

THE  DETERMINATION  OF  CHLORINE  IN  A 
SOLUBLE   CHLORIDE 

The  sample  may  be  an  artificial  mixture  of  the  chloride  and 
carbonate  of  sodium. 

Procedure.  Weigh  out  into  3oo-cc.  beakers,  two  portions, 
each  sufficient  to  contain  about  0.12  g.  of  sodium  chloride,  and 
treat  each  as  follows :  Dissolve  the  sample  in  50  cc.  of  water, 
run  in  from  a  burette  30.00  cc.  of  the  standard  silver  nitrate 
solution,  and  carefully  acidify  the  mixture  with  dilute  nitric 
acid.  Heat  to  boiling,  see  that  the  liquid  is  distinctly  acid,  and 
filter.  Receive  the  filtrate  and  washings  in  a  3oo-cc.  Erlenmeyer 
flask.  To  the  united  filtrate  and  washings,  which  should  have 
a  volume  of  about  150  cc.,  add  10  cc.  of  6-normal  nitric  acid  and 
5  cc.  of  the  indicator  solution,  and  titrate  with  the  standard 
thiocyanate  solution,  as  already  described.  Calculate  the  per- 
centage of  chlorine  in  the  sample. 

NOTES.  —  i.  Since  silver  chloride  is  several  times  as  soluble  as  silver 
thiocyanate,  the  former  must  be  filtered  off  before  the  titration  of  the 
excess  of  silver  nitrate ;  otherwise  the  silver  chloride  would  react  with  the 
thiocyanate  solution  and  render  the  end-point  uncertain.  This  behavior 
is  best  represented  by  the  following  system  of  equilibria : 


JC1-+ 

lAg+ 


(Solid)     (Diss'd)       lAg+  (+CNS-^AgCNS^±AgCNS). 

(Diss'd)       (Solid) 

That  is,  if  the  silver  chloride  were  left  in  the  mixture  during  the  titration, 
owing  to  the  slow  conversion  of  the  soluble  (colored)  thiocyanate  com- 
pounds into  insoluble  silver  thiocyanate,  there  would  be  no  permanent 
end-point. 

2.  Silver  bromide  and  silver  iodide  are  less  soluble  than  silver  thio- 
cyanate, so  that  in  the  determination  of  bromine  and  iodine  by  this  method 
it  is  not  necessary  to  filter. 

3.  Soluble  chlorates,  etc.,  may  be  determined  by  this  method  by  first 
reducing  them  to  the  corresponding  halides  (e.g.  with  sulphurous  acid), 
and  then  determining  the  latter. 

4.  For  other  uses  of  precipitation  methods,  see  Part  IV,  Problems,  90 
91,  92,  and  93. 


PART   IV 

STOICHIOMETRY 

Preliminary  Discussion.  The  stoichiometrical  problems  met 
with  in  analytical  work  are,  as  a  rule,  neither  hard  to  compre- 
hend nor  difficult  to  solve.  The  beginner  will  find  that  a  mod- 
erate amount  of  time  devoted  to  the  intelligent  study  of  these 
problems  will  enable  him  rapidly  to  make  the  calculations 
necessary  for  the  interpretation  of  analytical  data;  the  ability 
to  do  this  is  at  least  as  important  as  the  manipulative  skill  by 
which  the  data  are  obtained. 

It  cannot  be  too  strongly  emphasized  that,  in  making  such  cal- 
culationSj  the  beginner  should  from  the  outset  strive  to  take  the 
shortest  and  most  direct  route  to  the  result.  With  a  little  practice, 
the  student  who  is  not  unacquainted  with  the  reactions  of  an- 
alytical chemistry  should  soon  acquire  the  ability  to  recognize 
at  once,  upon  the  inspection  of  a  problem,  the  factors  which 
will  lead  most  directly  to  its  solution,1  as  well  as  the  equivalent 

1  Of  course  most  analytical  problems  can  be  solved  in  stages,  by  means  of  a 
series  of  proportions,  and  it  is  perhaps  only  natural  that  most  beginners  should 
have  a  predilection  for  this  method.  In  the  examples  given,  however,  the 
common  factors  have  been  eliminated,  and  the  problems  solved  hi  a  single  opera- 
tion. The  beginning  student  will  better  appreciate  the  advantages  of  the  shorter 
method  upon  comparing  the  solutions  given  of  problems  IV  and  V  with  the  follow- 
ing roundabout  method  of  arriving  at  the  same  results  : 

iv.  (a)  10  Fe :  2  KMnO4= 0.005  :  w. 

^  =  316/558X0.005  =  0.00283  g-  KMnO4  per  cc. 
(6)          2  KMnO4 :  5  H2C2O4= 0.00283  :  x. 

#=450/316X0.00283  =  0.00403  g.  H2C2O4  per  cc. 
(c)          5  H2C2O4 :  5  CaC2O4 = 0.00403  :  y. 

y  =  640/450X0.00403  =  0.005 73  g-  CaC2O4  per  cc. 
(rf)  5  CaC2O4 :  5  CaO = 0.005  73  :  *• 

2  =  280.5/640.5X0.00573=0.00251  g.  CaO  per  cc.    Ans. 


160     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

relationships  of  the  substances  involved.  To  do  this,  however, 
he  should  understand  and  bear  in  mind  the  relationships  and  dif- 
ferences which  exist  between  chemical  and  physical  units,  —  such 
as  atoms,  molecules,  and  equivalents  on  the  one  hand,  and  grams 
and  cubic  centimeters  on  the  other. 

Detailed  solutions  of  a  few  typical  problems  are  given  below. 
The  student  should  study  these  careiully  until  they  are  fully 
understood. 

i.  A  sample  of  a  soluble  chloride  weighing  0.200 7  g.  yields  on 
analysis  0.4920  g.  of  silver  chloride;  what  percentage  of  chlorine 
does  it  contain? 

{    From  the  proportion, 

Wt.  of  chlorine  in  sample :  Wt.  of  sample  =  %  of  chlorine :  100, 

...     ,    .        , ,    ,  Wt.  of  chlorine  in  sample  ~    r  , ,    . 

it  is  obvious  that ^—   - 100  =  %  of  chlorine. 

Wt.  of  sample 

Also,  since  the  chlorine  contained  in  the  sample  is  identical  with 
that  which  is  later  contained  in  tlie  silver  chloride  precipitate, 
we  have  the  proportion, 

Cl :  AgCl= Wt.  of  chlorine .  Wt.  of  silver  chloride, 

Cl 
or,  •  Wt.  of  silver  chloride =Wt.  of  chlorine. 

Substituting  this  value  in  the  preceding  equation,  we  get, 

-^-  •  Wt.  of  silver  chloride  354(5  -  0.4020 

AgC1    _      f = 100=  H^34 

Wt.  of  sample  0.2007 

=  60.50%  of  Cl. 

v.  (a)  AgCl:HCl=o.i527:*. 

x  =  36.46/143.34X0.1527  =  0.0388  g.  HC1  in  20.50  cc. 
(6)  20.50:  1000  =  0.0388 :  y. 

y=  1000/20.5X0.0388=  1.893  g-  HC1  in  one  liter. 
(c)  36.46:1.893  =  1:2. 

z=  1.893/36.46=0.0519  N.    Ans. 


STOICHIOMETRY  161 

A  chemical  factor  expresses  the  quantity  by  weight  of  an  ele- 
ment or  compound  which  is  equivalent  to  one  part  by  weight  of 
some  other  substance.  For  example,  the  ratio  or  factor 

Ag       107.88 

-T—  ^r=  -  '•  -  =  0.7526 

AgCl     143.34 

tells  us  that  one  gram  of  silver  chloride  contains  0.7526  g.  of 
silver,  and  if  we  wish  to  calculate  what  weight  of  silver  there  is 
in  a  specific  weight  of  silver  chloride,  we  simply  multiply  the 
latter  by  this  factor;  e.g.  10.15  g.  of  silver  chloride  contain 
10.15  Xo.7526  =  7.64  g.  of  silver. 

Again,  if  the  weight  of  FeO  which  corresponds  to  a  specific 
weight  of  Fe2C>3  is  desired,  the  factor  is 

2  FeO  _  143.68  _ 
Fe203"  159.68" 

And,  similarly,  if  it  is  wished  to  find  the  weight  of  K20  which 
corresponds  to  a  specific  weight  of  KC1,  the  factor  is 


In  the  calculation  of  these  (physical-unit)  factors,  the  equiva- 
lent relations  of  the  two  substances  must  be  kept  clearly  in 
mind  ;  thus  it  is  plainly  incorrect  to  express  the  ratio  of  potassium 

TT  O 

oxide  to  potassium  chloride  by  the  fraction  —  —  ,  since  each 

JxC/1 

molecule  of  K20  must  yield  upon  treatment  with  HC1  two  mole- 
cules of  KC1.  Similarly,  the  factor  for  the  conversion  of  Mn2P207 

into  Mn304  is  -  3    4    ;  for  two  molecules  of  MnsO^  contain- 
3  Mn2P2O7 

ing  six  atoms  of  manganese,  will  yield  three  molecules  of  Mn2P207, 
also  containing  six  atoms  of  manganese.  Carelessness  in  this 
respect  is  one  of  the  most  frequent  sources  of  error. 

ii.  How  many  cubic  centimeters  of  a  solution  containing  25  grams 
of  BaClz  •  2  HZ0  per  liter  will  be  required  to  precipitate  the  sulphur 
from  0.1073  Sram  °f  Pure  stibnite,  Sb^S^  as  BaSOt? 


162     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Each  molecule  of  80283  will  yield  upon  treatment  three  mole- 
cules of  H2S04,  and  these  will  require  three  molecules  of  barium 
chloride  for  precipitation.  We  therefore  arrive  at  the  propor- 
tion, 3  (BaCl2  •  2  H20)  :  80283=3;  :  0.1073,  where  x  represents  the 
weight  of  the  crystalline  salt  required.  That  is, 

_3(BaCl2.2H2Q) 

Sb2S3 

in  which  the  factor3  (Ba^2  g2  H20)  =2.177  indicates  the  quan- 

00203 

tity  by  weight  of  BaCl2  .  2  H20  which  is  required  to  precipitate 
the  sulphur  from  one  gram  of  Sb2S3  ;  this  factor  therefore  does  not 
differ  essentially  from  the  chemical  factors  previously  discussed. 
Finally,  since  each  cubic  centimeter  of  the  solution  contains 
0.025  g.  of  BaCl2  .  2  H20,  we  have, 


9-34  cc. 


0.025  0.025 


Hi.  What  volume  of  aqueous  ammonia  of  sp.  gr.  0.960,  contain- 
ing 9.91%  of  NH$,  will  be  required  to  precipitate,  as  Fe(OH)$,  the 
iron  contained  in  1.475  £•  of  Fe(NH4$Ot)2.6HzO?- 

Since  the  iron  is  to  be  precipitated,  after  oxidation,  as  Fe(OH)s, 
it  is  plain  that  each  atom  of  iron  will  require  three  molecules  of 
NH4OH,  which  in  turn  are  furnished  by  three  molecules  of  NH3. 


Therefore,        ,  -'I-^7=  wt.  of  NH3  required; 

Fe(NH4S04)2 


and,  since  each  cubic  centimeter  of  the  aqueous  ammonia  weighs 
0.960  g.,  and  contains  9.91%  of  NHs,  the  solution  contains 
0.960  Xo.0991  g.  of  NH3  per  cubic  centimeter.  That  is, 

51.10 

3O2.I6'1*475 

—  -  -    2.02  cc. 


0.960  Xo.0991  0.960  Xo.0991 


STOICHIOMETRY  163 

iv.  A  solution  of  potassium  permanganate  is  equivalent  to 
0.00500  g.  of  ferrous  iron  per  cubic  centimeter;  what  is  its  value 
in  terms  of  calcium  oxide  ? 

The  reactions  involved  in  the  volumetric  determinations  of 
iron  and  calcium  are  : 

10  FeSO4+2  KMnO4+9  H2SO4  =  2  KHS04+2  MnSO4 

+5Fe2(S04)3+8H20, 
and 

5  CaC2O4+2  KMn04+9  H2SO4  =  5  CaSO4+2  KHSO4 

+  2MnSO4+ioCO2+8H20; 

and  from  these  equations  it  is  clear  that,  in  this  case,  2  Fe++=c=CaO. 
That  is, 


CaO 


56.  i 
•  0.00500  =  —  -  •  0.00500  =  0.00251  g.  CaO. 


2  Fe  1 1 1. 6 

v.  If  20.50  cc.  of  hydrochloric  acid  yield  0.1527  g.  of  silver 
chloride,  what  is  the  normality  factor  of  the  solution? 

Although  silver  chloride  is  insoluble,  the  normality  factor  of 
the  acid  may  nevertheless  be  calculated  directly  from  the  weight 
of  the  precipitate  obtained.  One  liter  of  normal  hydrochloric 
acid,  containing  one  mol  of  HCl,  would  yield  143.34  g.  (i.e.  one 
mol)  of  silver  chloride,  whence  20.50  cc.  would  yield  0.14334 
X 20.50  g.  We  obtain,  therefore,  the  equation, 

o.: 


1      0.14334X20.50 

vi.  A  sample  of  stibnite  weighing  0.1793  g.  is  heated  with  strong 
HCl,  and  the  H2S  evolved  absorbed  by  means  of  sodium  hydroxide 
solution;  the  resulting  mixture  (containing  the  sulphur  as  Na^S) 
being  introduced  under  the  surface  of  a  solution  made  by  adding 
25  cc.  of  6-normal  HCl  and  50.00  cc.  of  o.n6o-normal  iodine 
to  500  cc.  of  water.  The  excess  of  iodine  is  titrated  with  0.0957- 


1  64     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

normal  sodium  thiosulphate  solution,  of  which  28.57  cc-  are  re- 
quired.   Calculate  the  percentage  of  (evolved)  sulphur  in  the  stibnite. 

(H2S+I2  =  2HI+S.) 
cc.      N.  F. 

50.00X0.1160  =  5.800  cc.  of  normal  iodine. 

28.57  Xo.OQ57  =  2.734  cc.  of  normal  thiosulphate. 
I.e.  the  H2S  required     3.066  cc.  of  normal  iodine. 

Since  normal  iodine  has  a  sulphur  value  of  °'°^20^=  0.01604  5  g- 
per  cubic  centimeter,  we  have, 

0.016035  X3.o66XioQ  Qf  ^ 


The  normality  factor  of  a  solution  expresses  the  value  of  the 
solution  per  cubic  centimeter  in  terms  of  a  normal  solution. 
For  example,  if  a  solution  is  known  to  be  one  half  normal  (i.e. 
N.  F.=  0.500),  it  is  obvious  that  i  cc.  of  it  is  equivalent  to 
1.000X0.500=0.500  cc.  of  a  normal  solution;  or  that  27.31  cc. 
of  it  are  equivalent  to  27.31X0.500  =  13.655  cc.  of  a  normal 
solution.  Knowing  the  normality  factors  of  a  series  of  solutions, 
therefore,  we  can  readily  reduce  the  different  volumes  of  the 
solutions  used  in  a  determination  to  a  common  standard  ;  and 
in  this  way  the  calculations  are  rendered  almost  as  simple  as 
if  the  tune  and  labor  had  been  expended  to  make  the  solutions 
all  of  exactly  the  same  strength,  say  one-tenth  normal. 

mi.  Indirect  methods  of  analysis  depend  upon  the  fact  that 
when  two  or  more  substances  are  made  to  undergo  the  same 
chemical  treatment  they  either  experience  a  relatively  different 
change  of  weight,  or  unit  weights  of  each  require  unequal  volumes 
of  a  standard  solution. 

For  example,  suppose  we  wish  to  determine  the  weight  of  NaCl 
and  of  KCl  in  a  mixture  of  the  two  salts.  The  mixture,  weighing 
a  grams,  may  be  converted  into  silver  chloride,  of  which  there  is 
formed,  say,  p  grams. 


STOICHIOMETRY  165 

Let  x  represent  the  weight  of  the  sodium  chloride,  and  y  that 
of  the  potassium  chloride,  and  we  have, 


and  AgCL       AgCl 

NaCl*hKCl  y    P 

If  we  designate  by  m  the  factor     ^      and  by  n  the  factor     % 
we  obtain, 


and  mx+ny=p, 

from  which  we  find  that 


or 


Indirect  analyses  may  in  general  be  calculated  by  means  of 
this  or  a  similar  general  equation. 
In  the  above  example, 


and  1^—^  =  0.5297. 

If  these  values  are  substituted  in  the  general  equation,  we  obtain, 
x  =  1.888  #-3.628  a. 

Consequently,  in  order  to  determine  the  weight  of  sodium 
chloride  in  the  mixed  sample  it  is  only  necessary  to  determine 
the  values  of  a  and  p,  multiply  them  by  3.628  and  1.888,  re- 
spectively and  subtract  the  first  product  from  the  second. 

The  sa  me  analysis  might  be  performed  by  weighing  the  mixed 
chlorides  in  a  platinum  crucible,  then  changing  them  to  sulphates 
(by  evaporation  with  H2SO4)  ,  and  again  weighing.  In  this  case  also, 


m—n        m—n 
21. 547  #-25.181  a. 


i66     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

In  the  first  case,  the  coefficients  are  relatively  small,  and 
consequently  good  results  might  be  expected,  since  the  experi- 
mental errors  made  in  the  determination  of  a  and  p  are  multiplied 
by  only  3.63  and  1.89,  respectively.  In  the  latter  case,  however, 
the  coefficients  are  very  large,  and  the  unavoidable  analytical 
errors  would  have  to  be  multiplied  enormously  in  the  calculation ; 
the  method  is  therefore  worthless. 

Although  some  indirect  methods  may  appear  simple  and 
attractive  on  paper,  they  frequently  lead  to  impossible  values 
in  practice ;  so  that  extreme  caution  should  be  exercised  regard- 
ing the  use  of  an  indirect  method.  In  general,  if  accurate  and 
reliable  results  are  desired,  indirect  methods  of  analysis  should  be 
avoided. 

PROBLEMS 

GRAVIMETRIC   ANALYSIS 

1.  Calculate  the  chemical  factors  for  the  following :   KC1  from  K2PtCl6 ; 
K20  from  K2PtCl6;    P  from  Mg2P2O7;    Fe3O4  from  Fe2O3;    MnO2  from 
Mn3O4. 

2.  What  weight  of  Mn304  corresponds  to  0.5785  g.  of  Mn2P2O7?     To 
0.4327  g.  of  MnSO4? 

3.  A  sample  of  an  impure  ammonium  salt  weighing  0.4988  g.  is  con- 
verted into  (NH4)2PtCl6,  and  this  upon  ignition  yields  0.3258  g.  of  platinum. 
Calculate  directly  from  the  weight  of  platinum  the  percentage  of   NH3 
in  the  sample. 

4.  A  sample  of  phosphorus  pentoxide  weighing  0.2018  g.  yields  0.3132  g. 
of  Mg2P20?.     Calculate  the  percentage  of  P2O5  in  the  sample. 

5.  What  weight  of  a  silver  nitrate  solution  known  to  contain  2.31% 
of  Ag  will  be  required  to  precipitate  the  chlorine  from  25.0  cc.  of  a  solution 
containing  25.0  g.  of  BaCl2 .  2  H2O  in  one  liter? 

6.  If  25.0  cc.  of  sodium  chloride  solution  yield  0.1434  g.  of  silver  chlo- 
ride, what  is  the  strength  of  the  solution  in  grams  of  the  salt  per  liter  ?     In 
mols  per  liter? 

7.  How  many  cubic  centimeters  of  a  solution  containing   25.0  g.  of 
BaCl2 .  2  H20  per  liter  will  be  required  to  precipitate,  as  BaSO4,  the  sul- 
phuric acid  formed  upon  oxidizing  0.2543  g.  of  FeS2  with  fuming  nitric  acid? 

8.  How  many  cubic  centimeters  of  hydrochloric  acid  of  sp.  gr.  1.050, 
containing  10.17%  of  HC1,  will  it  take  to  precipitate  the  silver  from  a  solu- 
tion containing  0.8430  of  silver  sulphate? 


STOICHIOMETRY  167 

9.  How  "many  cubic  centimeters  of  hydrochloric  acid  of  sp.  gr.  1.040, 
containing  8.16%  of  HC1,  will  be  required  to  dissolve  one  gram  of  calcium 
carbonate  ? 

10.  What  weight  of  Mn2P2O7  is  it  possible  to  prepare  from  50.0  cc.  of  a 
permanganate  solution  which  contains  4.500  g.  of  KMnO4  per  liter? 

11.  How  many  cubic  centimeters  of  a  solution  of  sp.  gr.  1.116,  con- 
taining 10.06%  of  NaOH-  will  it  take  to  neutralize  a  solution  containing 
5.00  g.  of  NaHS04?   5.00  g.  of  KHSO4? 

12.  A  sample  of  impure  potassium  sulphide  weighing  0.4320  g.  is  treated 
with  hydrochloric  acid,  and  by  means  of  ammoniacal  hydrogen  peroxide 
solution  the  hydrogen  sulphide  evolved  is  converted  into  ammonium  sul- 
phate.   This  yields  0.8034  g.  of  BaSO4.     Calculate  the  percentage  of  K2S 
in  the  sample. 

13.  A  sample  of  stibnite  weighing  1.078  g.,  upon  being  analyzed  by  the 
method  indicated  in  Problem  12,  yields  0.6750  g.  of  BaSO4.     Assuming 
the  sulphur  to  be  present  wholly  as  Sb2S3,  calculate  the  percentage  of  the 
latter  in  the  mineral. 

14.  How  many  cubic  centimeters  of  aqueous  ammonia  of  sp.  gr.  0.96, 
containing  9.91  %  of  NH3,  will  be  required  to  precipitate  the  aluminum  in 
0.8674  g.  of  KA1(SO4)2 .  12  H20  ?    How  many  cubic  centimeters  of  6-normal 
ammonia  ? 

15.  What  volume  of  the  ammonia  water  first  referred  to  in  Problem  14 
will  it  take  to  neutralize  10.0  cc.  of  hydrochloric  acid  of  sp.  gr.  1.12,  con- 
taining 23.81%  of  HC1?    To  neutralize  10.0  cc.  of  6-normal  hydrochloric 
acid? 

16.  If  15.0  cc.  of  a  solution  of  barium  chloride  yield,  upon  evaporation 
with  hydrochloric  acid  and  gentle  ignition,  1.563  g.  of  the  anhydrous  salt, 
what  is  the  strength  of  the  solution  in  mols  per  liter?    In  equivalents  per 
liter? 

17.  A  solution  contains  2.25%  by  weight  of  BaCl2 .    What  volume 
of  a  solution  containing  5.000  g.  of  Ag2S04  per  liter  will  be  required  to  pre- 
cipitate the  chlorine  in  5.000  g.  of  the  barium  chloride  solution  ?    What 
will  be  the  weight  of  the  precipitate  formed  ? 

18.  A  sample  of  pyrite  weighing  0.2500  g.  yields  upon  analysis  0.8020  g. 
of  BaS04.     Upon  the  assumption  that  the  sulphur  is  wholly  present  as 
FeS2,  calculate  the  percentage  of  the  latter  in  the  sample. 

19.  What  volume  of  bromine  water  containing  30  g.  of  bromine  per 
liter  will  be  required  to  oxidize  the  iron  in  1.75  g.  of  FeSO4  .  7  H2O? 

20.  What  volume  of  aqueous  ammonia  (sp.  gr.,  0.96,  containing  9.91% 
of  NH3)  will  be  required  to  precipitate  the  iron,  after  oxidation  with  hydro- 
gen peroxide,  from  a  solution  containing  0.750  g.  of  Fe(NH4S04)2 .  6  H2O 


168     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

and  12.0  cc.  of  hydrochloric  acid  (sp.  gr.,  1.12,  containing  23.8%  of  HC1)? 

21.  A  mixed  sample  of  CaO,  Ca(OH)2,  and  CaC03  weighing  0.5896  g. 
is  evaporated  with  excess  sulphuric  acid,  and  gently  ignited ;  the  residue 
weighs  0.8651  g.    What  volume  of  6-normal  hydrochloric  acid  will  be  re- 
quired to  convert  5.00  g.  of  the  sample  into  calcium  chloride? 

22.  The  ignited  precipitate  of  ferric  and  aluminum  oxides  from  1.497  g- 
of  a  mineral  weighs  0.4196  g. ;  after  ignition  in  a  current  of  hydrogen  the 
product  weighs  0.3311  g.,  the  ferric  oxide  being  reduced  to  metallic  iron. 
Calculate  the  percentage  of  Fe2C>3  and  of  AUOs  in  the  mineral. 

23.  A  sample  of  pyrolusite  weighing  0.5124  g.  is  heated  in  the  presence 
of  dilute  sulphuric  acid  with  an  excess  of  oxalic  acid,  and  the  gas  evolved 
is  absorbed  in  a  weighed  bulb  containing  potassium  hydroxide.     The  gain 
in  weight  of  the  bulb  is  found  to  be  0.4789  g.     Calculate  the  percentage  of 
Mn02  in  the  pyrolusite. 

24.  What  volume  of  o.5-normal  ammonium  oxalate  solution  will  be 
required  to  precipitate,  as  CaC2O4 .  H20,  the  calcium  from  one  gram  of 
apatite,  [(^(POOtJi  -  CaF2? 

25.  What  volume  of  a  solution  containing  66  g.  of  (NH4)2HP04  per  liter 
will  be  required  to  precipitate,  as  ZnNH4P04,  the  zinc  from  0.9786  g.  of  a 
brass  which  contains  30.15%  of  zinc?    What  is  the  normality  of  this  solu- 
tion as  a  precipitant  for  zinc  ? 

26.  How  many  grams  per  liter  of  K2Cr20:  must  a  solution  contain  in 
order  that,  by  reduction  of  the  chromium  (with  HC1  and  S02),  precipitar 
tion  with  ammonia,  and  ignition  of  the  precipitate  in  a  current  of  hydro- 
gen, a  50.0  cc.  portion  shall  yield  0.3752  g.  of  Cr208? 

27.  What  volume  of  a  solution  containing  25.0  g.  of  KH3(C204)2 .  2  H2O 
per  liter  will  be  required  to  precipitate  the  calcium  from  0.976  g.  of  a  lime- 
stone which  yields,  besides  a  small  quantity  of  an  insoluble  residue,  2.14% 
of  Fe20a,  9.56%  of  MgO,  and  45.36%  of  C02,  assuming  iron,  magnesium, 
and  calcium  to  be  present  wholly  as  carbonates  (the  iron  of  course  as  ferrous 
carbonate)  ? 

28.  What  volume  of  6-normal  sulphuric  acid  will  be  required  to  replace 
the  nitric  acid  in  the  salts  obtained  upon  evaporating  to  dryness  with  nitric 
acid  4.984  g.  of  brass,  if  the  brass  contains  65.98%  of  Cu,  31.42%  of  Zn, 
1.84%  of  Sn,  and  0.76%  of  Pb? 

29.  A  sample  of  silicate  mineral  weighing  1.0245  g.  yields  0.2602  g.  of 
potassium  and  sodium  chlorides;  and  the  mixed  chlorides  yield  0.4304  g. 
of  K2PtCl6.     Calculate  the  percentage  of  Na20  in  the  sample. 

30.  A  solution  of  chloroplatinic  acid  contains  0.050  g.  of  Pt  per  cubic 
centimeter.    What  is  the  minimum  volume  with  which  0.2602  g.  of  mixed 
sodium  and  potassium  chlorides  must  be  evaporated  in  order  to  insure  the 


STOICHIOMETRY  169 

complete  conversion  of  the  alkali  metals  into  chloroplatinates,  no  matter 
in  what  proportions  the  two  chlorides  may  exist  in  the  mixture  ? 

31.  A  sample  of  phosphate  rock  contains  0.87%  of  moisture  and  91.92% 
of  calcium  phosphate.    Calculate  the  percentage  of  Caa(P04)2  which  is 
present  on  the  dry  basis. 

32.  If  2.497  g.  of  a  fertilizer  containing  4.45%  of  moisture  yields  0.3150  g. 
of  Mg2P207j  what  is  the  percentage  of  P205  on  the  dry  basis? 

33.  Upon  treatment  with  sulphuric  acid,  1.430  g.  of  a  salt  yields  0.5952  g. 
of  Na2S04  and  101.5  cc-  of  COz,  measured  moist  at  17°  C.  and  757  mm. 
Calculate  the  percentages  of  Na20  and  C02  in  the  salt.    (Tension  of  aqueous 
vapor  at  17°=  14.45  mm.) 

34.  What  weight  of  water  is  present  in  one  liter  of  air  which  is  50% 
saturated  with  moisture  at  17°  and  748  mm.?     (See  problem  33.) 

35.  If  in  the  analysis  of  a  substance  an  error  of  0.1%  is  unavoidable, 
how  accurately  is  it  necessary  to  weigh  a  sample  of  200  mg.  ?    A  sample  of 
five  grams?    (See  p.  4.) 

36.  1.3250  g.  of  pure  Na2COs  is  dissolved  in  water  and  the  solution 
made  up  accurately  to  250.0  cc.    A  portion  is  carefully  transferred  without 
loss  to  a  platinum  dish  by  means  of  a  pipette  supposed  to  deliver  50.00  cc. 
of  liquid.    After  evaporation  with  hydrochloric  acid,  and  ignition,  the 
sodium  chloride  residue  is  found  to  weigh  0.2927  g.    What  volume  of  this 
solution  does  the  pipette  actually  deliver? 

37.  0.7500  g.  of  a  substance  containing  chlorine  and  bromine  yields 
0.5000  g.  of  Ag(Cl,  Br).    This  mixture  is  heated  in  a  current  of  chlorine, 
which  converts  the  bromide  of  silver  into  the  chloride,  and  the  loss  in  weight 
due  to  this  change  is  found  to  be  0.0683  g-     Calculate  the  percentages  of 
chlorine  and  bromine  in  the  sample. 

38.  From  the  following  data,  calculate  the  percentage  of  each  salt 
present  in  a  mixture  of  sodium  chloride,  bromide,  and  iodide :  Weight  of 
sample,  0.1500  g. ;  weight  of  precipitate  obtained  by  distilling  the  solution 
with  nitrous  acid  and  converting  the  iodine  into  silver  iodide,  0.1056  g. ; 
weight  of  silver  chloride  and  bromide  from  the  solution  after  removal  of 
the  iodine,  0.1784  g. ;  weight  of  this  precipitate  after  conversion  of  the 
whole  into  silver  chloride,  0.1623  g. 

39.  A  sample  of  baking  powder,  which  is  known  to  contain  only  NaHCOa 
and  KHC^-jOe,  in  equivalent  proportions,  and  starch,  yields  upon  treat- 
ment with  water  12.0%  by  weight  of  C02.     Calculate  the  percentage  of 
each  salt  in  the  sample. 

40.  A  salt  containing  barium,  chlorine,  and  water  of  hydration  gave 
upon  analysis  the  following  data:  Weight  of  sample,  i.oooo  g. ;  weight 
after  heating  (water  driven  off),  0.8522  g. ;  weight  of  silver  chloride  ob- 


170     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

tained,  i .  1 73  5  g. ;  weight  of  barium  sulphate  obtained,  0.9594  g.     Calculate : 
(a)  the  percentage  of  each  constituent ;  (6)  the  formula  of  the  compound. 

41.  If  two  i.oooo-g.  samples  of  a  substance  containing  10.00%  of  MgO 
are  weighed  out,  and  the  precipitate  of  MgNH4PO4 .  6  H20  is  in  one  case 
contaminated  with  0.0250  g.  of   Mg3(PO4)2,  and  in  the  other  case  with 
0.0250  g.  of  Mg[(NH4)2P04]2,  what  percentages  of  MgO  will  be  found  if 
the  calculations  are  based  upon  the  assumption  that  the  ignited  precipi- 
tate in  each  case  consists  entirely  of  Mg2P2O7? 

42.  The  carbonates  of  calcium,  strontium  and  barium  obtained  from  a 
lo-liter  sample  of  mineral  water  are  converted  into  the  anhydrous  nitrates, 
and  the  calcium  nitrate  is  extracted  with  absolute  alcohol-ether  mixture. 
The  residue  is  dissolved  in  water,  the  barium  separated  from  the  strontium, 
as  BaCr04,  and  the  strontium  precipitated  from  the  nitrate  with  sulphuric 
acid  and  alcohol.    There  are  finally  obtained  0.8507  g.  of  CaO,  0.1324  g. 
of  SrS04,  and  0.1072  g.  of  BaCrO4;  calculate  the  content  of  the  water  in 
milligrams  per  liter  (i.e.  parts  per  million)  of  Ca,  of  Sr,  and  of  Ba. 

43.  How  many  cubic  centimeters  of  sulphuric  acid  of  sp.  gr.  1.840,  con- 
taining 95.6%  of  H2SO4,  must  be  added  to  i  liter  of  sulphuric  acid  of  sp. 
gr.  1.560,  containing  65.1%  of  H2SO4,  to  obtain  a  solution  containing  75.0% 
of  H2SO4? 

44.  A  fuming  sulphuric  acid  contains  25.5%  of  non-hydrated  SO3.    How 
many  grams  of  98.2%  H2SO4  must  be  added  to  100  g.  of  the  fuming  acid  to 
give  a  product  containing  100%  of  H2SO4? 

45.  A  limestone  contains  90.0%  of  CaCO3,  3.50%  of  MgC03,  3.00%  of 
CaSO4-2  H20,  1.25%  of  FeCO3,  and  2.25%  of  anhydrous  siliceous  material. 
What  numerical  difference  would  you  expect  to  find  between  the  loss  on 
ignition  and  the  true  percentage  of  CO2  ? 

46.  An  ore  contains  28.15%  °f  nickel,  and  0.5000  g.  samples  are  taken 
for  analysis.     In  one  sample  the  element  is  determined  by  electrolysis,  as 
metallic  nickel,  while  in  a  second  sample  it  is  determined  by  means  of 
dimethylglyoxime,  as  Ni(C4H7N2O2)2.     If  the  algebraic  sum  of  the  errors 
involved  in  each  determination  were  equivalent  to  a  negative  error  of 
1.7  mg.  of  the  substance  finally  weighed,  how  much  greater  would  the 
percentage  error  be  in  the  first  determination  than  in  the  second? 

47.  An  electric  current  is  passed  simultaneously  through  a  series  of 
three  electrolytic  cells  which  contain  water  acidified  with  sulphuric  acid, 
an  ammoniacal  solution  of  nickel  sulphate,  and  molten  silver  chloride. 
What  is  deposited  upon  the  cathode  in  each  of  the  other  cells,  and  how 
many  grams,  in  the  time  in  which  one  liter  of  hydrogen,  measured  moist 
at  17°  and  746  mm.,  is  liberated  from  the  water?     (Tension  of  aqueous 
vapor  at  17°=  14.45  mm.) 


STOICHIOMETRY  171 

48.  From  the  following  data,  calculate  the  percentages  of  nickel  and 
cobalt  in  the  steel:  Weight  of  sample,  1.124  g. ;  weight   of   nickel   and 
cobalt  obtained  upon  electrolysis,  0.1246  g. ;  weight  of  nickel  dimethyl- 
glyoximine,  Ni  (C4H7N2O2)2,  obtained  from  the  electrolytic  deposit,  0.4382  g. 

49.  A  mass  of  platinum  weighs  12.145  g-  m  au"j  11.580  g.  in  water,  and 
11.115  g.  in  sulphuric  acid.    What  is  the  specific  gravity  of  the  platinum? 
Of  the  sulphuric  acid? 

50.  A  quantity  of  pure  metallic  silver  weighing  1.0788  g.  is  dissolved  in 
nitric  acid  and  the  solution  made  up  to  the  mark  in  a  measuring  flask  gradu- 
ated to  contain  100.0  cc.    Three  portions  are  carefully  transferred  without 
loss  to  three  separate  beakers  by  means  of  a  pipette  known  to  deliver  25.00  cc. 
If  the  solution  remaining  in  the  flask,  together  with  the  liquid  finally  washed 
from  the  pipette,  yields  on  analysis  0.3560  g.  of  AgCl,  what  volume  of  liquid 
does  the  flask  actually  contain? 

VOLUMETRIC  ANALYSIS 

51.  If  25.00  cc.  of  hydrochloric  acid  yield  0.1435  g.  of  AgCl,  what  is  the 
normality  of  the  solution  ? 

52.  If  a  25.00  cc.  portion  of  acid  requires  21.50  cc.  of  0.526  N  alkali  for 
neutralization,  what  is  the  normality  of  the  acid?     Supposing  the  acid  to 
be  HC1,  what  weight  of  silver  chloride  will  10.00  cc.  of  it  yield  with  silver 
nitrate  ? 

53.  If  a  2.453  g.  sample  of  pure  anhydrous  sodium  carbonate  requires 
45.72  cc.  of  an  acid  for  neutralization,  and  if  41.90  cc.  of  the  acid  requires 
44.35  cc.  of  an  alkali,  what  is  the  normality  factor  of  each  solution? 

54.  A  sample  of  pure  calcite,  CaCOs,  weighing  2.150  g.  is  dissolved  in 
50.00  cc.  of  an  acid,  and  the  excess  of  acid  is  neutralized  with  29.12  cc.  of 
an  alkali  of  which  28.40  cc.  require  7.10  cc.  of  the  acid  for  neutralization. 
To  what  volume  must  one  liter  of  the  acid  be  diluted  in  order  to  make  it 
exactly  normal  ? 

55.  How  many  cubic  centimeters  of  0.526  N  acid  will  it  take  to  neu- 
tralize the  ammonia  set  free  upon  distilling  1.0378  g.  of  MgNH4PO4 .  6  H2O 
with  an  excess  of  caustic  alkali? 

56.  If  15.25  cc.  of  alkali  will  neutralize  20.00  cc.  of  a  solution  containing 
6.000  g.  of  KH3(C2O4)2  •  2  H2O  in  250.0  cc.,  what  is  the  normality  factor 
of  the  alkali? 

57.  In  the  analysis  of  a  feeding  stuff,  a  Kjeldahl  determination  is  carried 
out  with  a  sample  weighing  1.500  g.    The  ammonia  is  received  in  25.00  cc. 
of  0.500  N  acid,  and  the  excess  of  acid  is  found  to  require  12.50  cc.  of  a 
standard  alkali,  of  which  21.20  cc.  will  neutralize  18.00  cc.  of  the  acid. 
Calculate  the  percentage  of  nitrogen  in  the  sample. 


172     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

58.  What  weight  of  crude  cream  of  tartar  must  be  taken  for  titration 
in  order  that  twice  the  number  of  cubic  centimeters  of  0.2000  N  alkali 
required  may  numerically  equal  the  percentage  content  of  KH^H^Oe? 

59.  A  sample  of  soda  ash  weighing  25.05  g.  is  dissolved  and  made  up  to 
250.0  cc.,  and  one  fifth  of  this  solution  is  taken  for  titration.    What  must 
be  the  normality  of  the  standard  acid  (assuming  the  alkalinity  to  be  due 
only  to  Na2C03)  in  order  that  the  burette  reading  multiplied  by  two  may 
indicate  the  percentage  of  Na2C03  in  the  sample? 

60.  A  sample  of  caustic  soda  weighing  4.000  g.  is  dissolved  in  water  and 
made  up  to  one  liter.    A  100.0  cc.  portion  of  this  solution  requires  for 
neutralization  47.50  cc.  of  0.2000  N  acid.     A  second   100.0  cc.  portion, 
after  treatment  with  barium  chloride  in  slight  excess,  is  diluted  to  200.0  cc. 
and  allowed  to  settle,  and  50.0  cc.  of  the  clear  solution  require  11.50  cc.  of 
the  acid.     Calculate  the  percentages  of  NaOH  and  Na2C03  in  the  sample. 
(Neglect  the  volume  occupied  by  the  solid  precipitate.) 

61.  A  sample  of  Solvay  soda  weighing  3.750  g.  is  dissolved  in  water  and 
made  up  to  one  liter.    A  100.0  cc.  portion  of  this  solution,  titrated  in  the 
cold  with  o.i  ooo  N  acid,  with  the  use  of  phenolphthalein,  is  found  to  re- 
quire 29.95  cc-  of  the  acid;  the  burette  is  then  refilled  and  the  titration 
completed  at  the  boiling  temperature  of  the  solution,  35.15  cc.  more  of 
the  acid  being  required.     Calculate  the  percentages  of  Na2C03  and  NaHC03 
in  the  sample.     (Under  suitable  experimental  conditions,  phenolphthalein 
becomes  colorless  in  the  cold  as  soon  as  the  carbonate  has  been  wholly 
converted  into  bicarbonate.) 

62.  A  sample  of  sirupy  phosphoric  acid  weighing  5.767  g.  is  dissolved  in 
water  and  made  up  to  one  liter.    A  100.0  cc.  portion  of  the  solution  is  treated 
with  sodium  acetate  and  silver  nitrate  in  excess,  whereby  the  phosphate  is  quan- 
titatively precipitated  as  AggPC^.    Phenolphthalein  is  added  to  the  filtrate 
and  washings,  and  the  solution  titrated  with  0.500  N  alkali,  of  which  27.25  cc. 
are  required.     Calculate  the  percentage  of  H3P04  in  the  original  sample. 

63.  A  sample  of  Chili  saltpeter  weighing  1.025  g.  is  treated  in  sodium 
hydroxide  solution  with  pulverized  Devarda's  alloy  (50%  Cu,  45%  Al, 
5%  Zn),  which  reduces  the  nitrogen  to  ammonia;  the  ammonia  is  distilled 
into  25.00  cc.  of  0.463  ^V  acid,  and  the  excess  of  acid  requires  5.01  cc.  of 
0.212  N  alkali  for  neutralization.     Assuming   that  nitrogen  was  wholly 
present  as  NaN03,  calculate  the  percentage  of  the  latter  in  the  sample. 

64.  A  sample  of  strontium  nitrate  weighing  10.53  g-  is  dissolved  in  water 
and  made  up  to  one  liter.    One  tenth  of  this  solution  is  distilled,  in  the 
presence  of  alkali,  with  an  excess  of  titanous  chloride,  which  reduces  the 
nitrate  to  ammonia  (KN03+8  Ti(OH)3+6  H20=KOH+8  Ti(OH)4+NH3). 
The  ammonia  is  received  in  25.00  cc.  of  0.500  N  acid,  and  the  excess  of  acid 


STOICHIOMETRY  173 

requires  10.16  cc.  of  0.250  N  alkali.    Assuming  that  the  nitrogen  was  wholly 
present  as  Sr(N03)2,  calculate  the  percentage  of  the  latter  in  the  sample. 

65.  How  many  grams  of  KH3(C204)2 .  2  H20  will  it  take  to  prepare  one 
liter  of  a  0.500  N  solution,  to  be  used  as  a  standard  acid?    How  many  to 
prepare  one  liter  of  a  o.iooo  N  solution,  to  be  used  as  a  reducing  agent  in 
connection  with  potassium  permanganate  ? 

66.  How  many  grams  of  K2Cr207  per  liter  will  be  required  to  prepare  a 
solution  of  such  strength  that  each  cubic  centimeter  shall  indicate  2.00% 
of  iron,  when  a  sample  weighing  0.2792  g.  is  used  for  analysis?    What 
is  the  normality  factor  of  this  solution  ? 

67.  From  the  following  data,  calculate  the  percentage  of  iron  in  the 
ore:  Weight  of  sample,  0.2186  g. ;  the  reduced  iron  solution  requires  for 
oxidation  25.14  cc.  of  0.0996  N  permanganate  solution. 

68.  What  is  the  maximum  weight  of  an  ore  containing  70.00%  of  iron 
which  can  be  taken  for  analysis  without  having  to  refill  a  3o-cc.  burette, 
if  the  permanganate  solution  is  0.1025  iV? 

69.  A  calcium  oxalate  precipitate,  obtained  from  0.8432  g.  of  a  rock, 
is  decomposed  with  dilute  sulphuric  acid  and  made  up  to  250.0  cc. ;  of  this 
a  50.0  cc.  portion  is  titrated  with  0.1012  N  permanganate  solution.    If 
27.35  cc-  of  the  latter  are  required,  what  percentage  of  CaO  does  the  rock 
contain? 

70.  11.56  cc.  of    nitric  acid  of  sp.  gr.  1.19  are  diluted  to   250.0  cc. ; 
20.00  cc.  of  this  solution  are  found  to  require  12.92  cc.  of  0.410  N  alkali. 
Calculate  the  percentage  of  HN03  in  the  acid  of  sp.  gr.  1.19. 

71.  The  Sb2S3  precipitate  obtained  from  the  solution  of  an  ore  is  dis- 
solved in  sodium  sulphide  solution,  and  this  is  evaporated  and  fumed  with 
an  excess  of  sulphuric  acid.    The  residue  is  then  dissolved  in  dilute  hydro- 
chloric acid,  and  the  antimony  oxidized  from  the  trivalent  to  the  penta- 
valent  condition  by  means  of  a  standard  solution  of  permanganate.    The 
sample  of  ore  weighs  0.2749  g.  and   24.17  cc.  of  0.1025  N  permanganate 
solution  are  used  in  the  titration ;  what  is  the  percentage  of  antimony? 

72.  What  is  the  normality  factor  of  an  acid,  of  which  25.37  cc-  are  equiva- 
lent to  1.263  g.  of  KN03  when  the  nitrogen  of  the  latter  is  reduced  in  alkaline 
solution  to  ammonia  and  this  is  distilled  off.  and  received  in  the  acid  solution? 

73.  From  the  following  data,  calculate  the  percentage  of  Mn02  in  the 
ore:  Weight  of  sample,  0.2000  g. ;  after  heating  this  in  the  presence  of 
sulphuric  acid  with  50.00  cc.  of  o.iooo  N  oxalic  acid,  the  excess  of  oxalic 
acid  requires  8.50  cc.  of  o.iooo  N  permanganate  solution. 

74.  From  the  following  data,  calculate  the  percentage  of  Mn02  in  the 
ore:  Weight  of  sample,  0.2400  g. ;  this  is  boiled  with  hydrochloric  acid 
and  the  distillate  received  in  an  excess  of  potassium  iodide  solution, 


174     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

the  liberated  iodine  requiring  25.51  cc.  of  0.2000  ^V  sodium  thiosulphate 
solution. 

75.  A  sample  of  mineral  substance  weighing  i.ooo  g.  is  taken  for  analysis. 
In  the  determination  of  the  iron,  the  ferric  solution  is  completely  reduced 
by  means  of  sulphurous  acid,  and  the  excess  of  the  latter  removed  by  the 
passage  of  carbon  dioxide  through  the  boiling  solution;  the  iron  then  re- 
quires 28.17  cc-  °f  o.iooo  N  permanganate  solution.     Calculate  the  per- 
centage of  iron  in  the  substance. 

76.  What  weight  of  pyrolusite  containing  86.50%  of  MnO2  will  oxidize 
the  same  amount  of  oxalic  acid  as  50.0  cc.  of  a  permanganate  solution,  if 
10.00  cc.  of  the  latter  will  liberate  0.1905  g.  of  iodine  from  an  acidified  solu- 
tion of  potassium  iodide  ? 

77.  What  weight  of  Fe(NH4SO4)2 .  6  H2O  will  reduce   50.0  cc.  of  a 
permanganate  solution,  of  which  10.00  cc.  will  liberate  from  an  acidified 
solution  of  potassium  iodide  a  quantity  of  iodine  sufficient  to  react  with 
15.25  cc.  of  0.1025  N  thiosulphate  solution? 

78.  A  standard  solution  of  permanganate  will  oxidize  0.00730  g.  of  ferrous 
iron  per  cubic  centimeter ;  what  is  the  value  of  the  same  solution  in  terms  of 
(a)  H2C2O4 .  2  H20 ;  (6)  KNO2 ;  (c)  H2O2 ;  (<J)  K4Fe(CN) « .  3  H2O ;  (e)  Mn  ? 

79.  The  calcium  oxalate  precipitate  obtained  from  0.2500  g.  of  calcite 
is  dissolved  in  an  excess  of  sulphuric  acid,  and  the  hot  solution  titrated  with 
a  solution  of  permanganate  of  which  each  cubic  centimeter  represents 
0.00735  g-  °f  Na2Cs04.     If  45.57  cc.  of  the  permanganate  solution  are  re- 
quired, what  percentage  of  calcium  does  the  mineral  contain  ? 

80.  If  0.1340  g.  of  sodium  oxalate  require  19.23  cc.  of  a  permanganate 
solution,  how  many  milligrams  of  ferrous  iron  will  each  cubic  centimeter 
of  the  permanganate  solution  indicate  ? 

81.  From  the  following  data,  calculate  the  percentage  of  manganese 
in  the  ore:  Weight  of  sample,  0.5027  g. ;  volume  of  permanganate  solu- 
tion required  to  oxidize  the  manganese,  36.60  cc. ;  value  of  the  perman- 
ganate solution  for  use  with  oxalic  acid,  0.0997  N- 

82.  2.400  liters  of  dry  air  (at  o°,  760  mm.)  and  50.00  cc.  of  o.oioo  N 
barium  hydroxide  solution  are  shaken  together,  and  the  excess  of  alkali 
is  found  to  require  35.06  cc.  of  o.oioo  ^V  acid.    What  volume  of  carbon 
dioxide  is  contained  in  10,000  volumes  of  the  dry  air? 

83.  If  a  permanganate  solution  is  equivalent  to  5.84  mg.  of  iron  per  cubic 
centimeter,  what  is  the  value  of  the  solution  in  terms  of  K4Fe(CN)6 .  3  H2O? 
In  terms  of  N20s? 

84.  What  weight  of  iodine  per  cubic  centimeter  will  be  liberated  by  a 
permanganate  solution  from  an  excess  of  hydriodic  acid,  if  the  perman- 
ganate solution  has  an  iron  value  of  4.98  mg.  per  cubic  centimeter? 


STOICHIOMETRY  175 

85.  A  sample  of  iron  wire  weighing  0.1408  g.,  and  containing  99.8% 
of  iron,  is  converted  into  ferrous  chloride  and  titrated  according  to  the 
Zimmermann-Reinhardt  method;  it  requires  25.15  cc.  of  a  permanganate 
solution.     What  weight  of  ore  must  be  taken  for  analysis  by  the  same 
method,  in  order  that  each  cubic  centimeter  of  the  permanganate  solution 
may  indicate  2.50%  of  iron? 

86.  25.00  cc.  of  a  certain  acid  are  found  to  require  23.67  cc.  of  an  alkali. 
If  28.15  cc-  of  the  acid  are  used  to  dissolve  0.5260  g.  of  pure  calcium  car- 
bonate, and  6. 6 1  cc.  of  the  alkali  are  required  to  neutralize  the  resulting 
solution,  what  is  the  normality  factor  of  the  acid?    What  is  that  of  the 
alkali? 

87.  The  calcium  oxalate  precipitate  from  0.5005  g.  of  a  mineral  requires, 
after  decomposition  with  sulphuric  acid,  43.06  cc.  of  a  permanganate  solu- 
tion which  has  a  Na2C2O4  value  of  6.70  mg.  per  cubic  centimeter.     Cal- 
culate the  percentage  of  CaO  in  the  mineral. 

88.  In  the  standardization  of  a  dichromate  solution,  a  sample  of  pure 
iron  weighing  0.2000  g.  is  converted  into  200  cc.  of  ferrous  chloride  solution, 
and  titrated.     In  the  subsequent  analysis  of  an  ore  with  the  dichromate 
solution,  an  equal  quantity  of  ferrous  iron  is  present,  but  the  solution  to  be 
titrated  has  a  volume  of  600  cc.     If  the  indicator  used  permits  the  recog- 
nition of  one  part  by  weight  of  ferrous  iron  in  100,000  of  solution,  what  error 
results  from  the  fact  that  the  two  titrations  are  made  at  different  volumes  ? 
(Assume  that  the  dichromate  solution  was  found  to  be  o.iooo  N,  and  that 
the  solutions  have  the  specific  gravity  of  water.) 

89.  A  sample  of  crystalline  ammonium  acetate  weighing  2.021  g.  is 
dissolved  in  water  and  made  up  to  200.0  cc.     One  half  of  this  solution  is 
distilled  with  an  excess  of  lime,  and  the  distillate  received  in  25.00  cc.  of 
0.500  N  acid;   the  second  half  is  distilled  with  phosphoric  acid  in  excess, 
and  the  distillate  received  in  45.00  cc.  of  0.500  N  alkali.    In  the  first  case, 
with  methyl  orange,  the  excess  of  acid  requires  10.20  cc.  of  the  standard 
alkali,  and  in  the  second  case,  with  phenolphthalein,  the  excess  of  alkali 
requires  15.52  cc.  of  the  standard  acid.     Calculate  the  formula  of  the  salt. 

90.  The  aqueous  solution  of  0.1361  g.  of  a  mixture  containing  only  sodium 
chloride  and  bromide  is  treated,  in  the  presence  of  nitric  acid,  with  25.00  cc. 
of  o.iooo  N  AgN03  solution,  and  the  precipitate  is  filtered  off  and  washed. 
The  filtrate  and  washings  require  9.70  cc.  of  o.iooo  N  thiocyanate  solution. 
Calculate  the  percentages  of  sodium  chloride  and  bromide  in  the  sample. 

91.  A  mixture  containing  soluble  chlorides  and  iodides  weighs  0.4500  g. 
This  is  treated,  in  the  presence  of  nitric  acid,  with  35.00  cc.  of  o.iooo  N 
silver  nitrate  solution,  and  the  precipitate  is  found  to  weigh  0.5000  g.     The 
excess  of  silver  nitrate  in  the  filtrate  and  washings  requires  11.10  cc.  of 


176     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

0.0500  N  thiocyanate  solution.      Calculate  the  percentages  of   chlorine 
and  iodine  hi  the  original  mixture. 

92.  To  20.00  cc.  of  a  solution  of  prussic  acid,  an  excess  of  sodium  hy- 
droxide and  a  very  little  potassium  iodide  are  added ;  this  solution  is  titrated 
with  o.iooo  N  silver  nitrate,  of  which  48.73  cc.  are  required  to  produce  a 
fahit  permanent  turbidity  of  silver  iodide.     Calculate  the  percentage  of 
HCN  in  the  original  solution,  assuming  its  specific  gravity  to  be  equal  to 
that  of  water. 

93.  A  mixture  containing  ootassium  cyanide  and  chloride,  and  weigh- 
ing 0.2037  g.,  is  dissolved  in  water  and  titrated  with  o.iooo  N  silver  nitrate 
solution,  of  which  14.41  cc.  are  required  to  produce  a  faint  permanent 
turbidity.     16.19  cc-  more  of  the  silver  nitrate  solution  are  added,  the  solu- 
tion is  slightly  acidified  with  nitric  acid,  and  the  filtrate  and  washings  from 
the  precipitate  require  for  titration  14.51  cc.  of  o.iooo   N   thiocyanate 
solution.     Calculate  the  percentages  of  KCN  and  KC1  in  the  sample. 

94.  A  solution  of  potassium  permanganate  is  equivalent  to  5.00  mg.  of 
iron  per  cubic  centimeter.    To  40.0  cc.  of  this  solution,  acidified  with  an 
excess  of  very  dilute  sulphuric  acid,  an  excess  of  potassium  iodide  is  added, 
and  the  liberated  iodine  is  titrated  with  a  solution  of  sodium  thiosulphate, 
of  which  34.85  cc.  are  required.     Calculate  the  normality  factor  of  the 
thiosulphate  solution. 

95.  A  sample  of  soda  weighing  4.973  g.  is  dissolved  in  water  and  made 
up  to  one  liter.    A  100.0  cc.  portion  of  this  solution,  carefully  titrated  in 
the  cold  with  phenolphthalein  as  an  indicator,  requires  48.90  cc.  of  0.0998  N 
acid.    A  second  portion  of  100.0  cc.  is  titrated  with  0.499  N  acid,  of  which 
13.81  cc.  are  required  with  methyl  orange  as  an  indicator.     Calculate  the 
percentages  of  Na?C03  and  NaOH  in  the  sample.     (Cf.  Problem  61.) 

96.  A  sample  of  bleaching  powder  weighing  7.092  g.  is  triturated  with 
water  and  made  up  to  one  liter.    A  50.0  cc.  portion  of  this  suspension, 
when  titrated  with  o.iooo  N  arsenious  acid,  with  potassium  iodide-starch 
paper  as  an  outside  indicator,  is  found  to  require  26.15  cc.  of  the  standard 
solution.    What  is  the  percentage  of  available  chlorine  in  the  bleaching 
powder? 

97.  If  43.60  cc.  of  a  thiosulphate  solution  require  40.15  cc.  of  an  iodine 
solution,  and  if  0.2118  g.  of  As20a  require  42.40  cc.  of  the  iodine  solution, 
what  is  the  normality  factor  of  each  solution  ? 

98.  A  sample  of  titaniferous  ore  weighing  0.3805  g.  is  fused  with  a  mix- 
ture of  K2S207  and  KF,  the  melt  dissolved  in  HC1,  and  the  iron  and  titanium 
reduced  with  zinc  in  an  atmosphere  of  hydrogen.    The  solution  is  then 
titrated  in  an  atmosphere  of  carbon  dioxide,  in  the  presence  of  i  g.  of  KSCN 
as  the  indicator,  with  o.iooo  N  ferric  chloride  solution,  of  which  19.34  cc. 


STOICHIOMETRY  177 

are  required.  Calculate  the  percentage  of  Ti02  in  the  ore.  (TiCl3-f  FeCla 
=TiCl4+FeCl2.) 

99.  In  50.0  cc.  of  a  solution,  containing  both  ferrous  and  ferric  sulphates, 
the  ferrous  iron  is  titrated  in  the  presence  of  sulphuric  acid  with  o.iooo  N 
permanganate  solution,  after  which  the  oxidized  solution  is  titrated  for  total 
iron  with  0.0997  N  titanous  chloride  solution,  with  potassium  thiocyanate 
as  the  indicator.  If  hi  the  first  titration  27.17  cc.  of  the  permanganate 
solution  are  used,  and  in  the  second  titration  46.98  cc.  of  the  titanous 
chloride  solution,  what  is  the  content  of  the  original  solution  in  grams  per 
liter  of  ferrous  and  of  ferric  iron? 

i  oo.  The  ammonium  phosphomolybdate  precipitate  obtained  from 
2.000  g.  of  steel  is  dissolved  in  dilute  aqueous  ammonia,  the  solution  acidified 
with  sulphuric  acid,  and  the  molybdenum  reduced  to  the  trivalent  con- 
dition by  passing  the  acid  solution  through  a  Jones'  reductor,  —  the 
reduced  solution  being  caused  to  enter  the  receiving  vessel  under  the  sur- 
face of  a  solution  of  ferric  sulphate  (2  Mo03+6  H=Mo203+3  H20;  and 
Mo2O3  +  3Fe2O3  =  2MoO3H-6FeO).  The  resulting  solution  is  at  once 
titrated  with  a  standard  solution  of  permanganate,  of  which  18.75  cc.  are  re- 
quired. If  the  permanganate  solution  has  an  iron  value  of  0.00540  g.  per 
cubic  centimeter,  what  is  the  percentage  of  phosphorus  in  the  steel? 


PART   V 

QUESTIONS 

Exercises  with  the  Balance. 

1.  What  is  the  purpose  of  weighing  ? 

2.  Explain  the  mechanical  theory  of  the  balance. 

3.  Give  five  conditions  which  must  be  satisfied  by  a  good  balance. 

4.  What  conditions  must  be  fulfilled  in  order  that  a  balance  may  1 
considered  properly  adjusted  for  use  ? 

5.  How  may  the  zero-point  of  a  balance  be  determined  ?    Illustrate. 

6.  Explain  the  following  methods  of  weighing:   (a)  Ordinary  methoi 
(6)  Weighing  by  double  vibrations ;  (c)  Method  of  Gauss ;  (d)  Method 
Borda. 

7.  Describe  a  procedure  for  the  calibration  of  a  set  of  weights. 

8.  Discuss^the  errors  in  weighing  which  may  be  due:  (a)  to  inequaliti 
in  length  in  the  beam  arms ;  (6)  to  the  buoyancy  of  the  atmosphere. 

9.  What  is  a  desiccator?    Explain  why  it  is  necessary,  and  give  tl 
principles  upon  which  its  use  is  based. 

The  Determination  of  Chlorine. 

1.  What  substances  would,  if  present,  interfere  with  this  determination! 

2.  Why  should  the  solution  be  acidified  with  nitric  acid?    Why  shou 
a  large  excess  of  nitric  acid  be  avoided  ? 

3.  Why  should  the  solution  not  be  heated  until  after  the  addition  of  tl 
silver  nitrate  ?    Why  is  it  then  heated  ? 

4.  What  are  the  advantages  of  washing  by  decantation  ?    In  washing 
precipitate,  whether  by  decantation  or  otherwise,  why  should  the  liquid  eac 
time  be  removed  as  far  as  possible  before  the  addition  of  fresh  wash  liquic 

5.  How  can  you  tell  when  the  precipitate  has  been  sufficiently  washed! 

6.  Why  is  the  filter  paper  ignited  separately  from  the  bulk  of  the  pr 
cipitate?    What  is  the  object  of  the  treatment  with  nitric  and  with  hydr< 
chloric  acid?    Explain. 

7.  What  is  the  effect  of  light  upon  silver  chloride?    Is  the  action  < 
diffused  daylight  a  serious  source  of  error? 

178 


QUESTIONS  179 

8.  Why  should  the  precipitate  be  heated  until  it  just  begins  to  fuse? 
What  is  the  effect  of  overheating  ?    Of  underheating  ? 

9.  Water  saturated  with  silver  chloride  at  100°  contains  about  22  mg. 
of  the  salt  per  liter.    Explain  why  the  precipitate  may  be  thoroughly  washed 
with  hot  water  without  undue  loss.    What  is  the  solubility  of  silver  chloride 
in  water  at  the  ordinary  room  temperature  ?    In  the  precipitation  of  chloride 
ion,  why  should  silver  nitrate  be  added  in  moderate  excess? 

10.  Given  an  aqueous  solution  of  silver  chloride  in  equilibrium  with  a 
quantity  of  the  solid  salt :  What  would  happen  upon  the  addition  of  (a)  a 
few  drops  of  silver  nitrate  solution?     (6)  a  few  drops  of  sodium  chloride 
solution?   (c)  a  large  excess  of  sodium  chloride?    Explain  in  each  case. 

11.  How  may  the  crucible  safely  and  readily  be  cleaned  after  the  ignition 
of  the  silver  chloride?    Write  equations  to  show  the  reactions  involved. 

12.  (a)  Explain  the  solubility  of  silver  chloride  in  each  of  the  following 
substances:  aqueous  ammonia;  potassium  cyanide  solution;  sodium  thio- 
sulphate  solution.     (5)  Is  silver  iodide  soluble  in  aqueous  ammonia?    In 
a  solution  of  potassium  cyanide  ?    Explain  your  answers. 

13.  What  other  substances  may  be  determined  in  a  similar  manner  in 
the  form  of  insoluble  silver  salts  ? 

14.  Why  are  Gooch  crucibles  preferable  to  ordinary  paper  filters;  es- 
pecially, for  example,  in  the  determination  of  iodides,  cyanides,  etc.  ? 

15.  Starting  with  the  native  mineral,  describe  the  treatment  which 
renders  the  asbestos  suitable  for  use  in  the  preparation  of  Gooch  crucibles. 

The  Determination  of  Iron  and  of  Sulphur  in  a  Soluble  Sulphate  of  Iron. 

Iron.  —  i .  A  mixture  consists  of  ferric  sulphate,  sodium  carbonate,  and 
potassium  sulphate,  each  of  which  is  soluble  in  cold  water.  Will  the  mix- 
ture dissolve  in  water  ?  Illustrate  your  answer  by  means  of  equations. 

2.  How  is  the  sample  taken  into  solution  for  this  analysis? 

3.  In  the  precipitation  with  ammonium  hydroxide,  why  must  the  iron 
be  present  wholly  in  the  ferric  condition? 

4.  How  is  the  ferrous  iron  oxidized  in  this  analysis?    Write  the  equa- 
tion.   How  can  you  tell  when  the  oxidation  is  complete  ? 

5.  How  may  it  be  ascertained  whether  the  original  sample  contains 
ferrous  iron?    Is  it  worth  while  to  make  this  test ?    Why? 

6.  What  advantage  is  to  be  gained  by  adding  the  ferric  salt  solution  to 
an  excess  of  aqueous  ammonia,  rather  than  the  ammonia  to  the  ferric  salt 
solution  ?    Explain. 

7.  How  is  the  first  precipitate  of  ferric  hydroxide  treated,  and  why? 
Explain  in  full. 

8.  Is  it  necessary  to  completely  wash  out  the  ammonium  chloride 
before  the  ferric  hydroxide  is  ignited?    Why? 


l8o     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

9.  What  precautions  are  to  be  observed  in  igniting  the  precipitate  of 
ferric  hydroxide? 

10.  Name  two  other  metals  which  may  be  determined  by  a  similar  pro- 
cedure.   What  additional  precautions  should  be  taken  in  their  determina- 
tion, and  why  ? 

11.  In  the  determination  of  these  three  metals,  would  it  be  equally  well 
to  use  sodium  hydroxide  as  the  precipitating  reagent  ?     Give  the  reason  for 
your  answer  in  the  case  of  each  metal. 

12.  What  is  the  effect  of  tartaric  or  citric  acid,  sugars,  etc.  upon  the  pre- 
cipitation of  ferric  hydroxide  by  means  of  ammonium  or  sodium  hydroxide  ? 
Do  any  other  metallic  ions  behave  like  Fe+++-ion  in  this  respect  ?    Explain. 

Sulphur.  —  i.   Why  must  nitrates  be  removed  before  the  precipitation 
with  barium  chloride  ?    How  is  this  done  ?    Write  the  reaction. 

2.  Name  some  other  substances  which,  if  present,  should  be  removed 
before  precipitation  with  barium  chloride. 

3.  Why  is  barium  chloride  chosen  as  the  reagent  for  sulphate,  rather 
than  lead  nitrate  or  strontium  chloride? 

4.  How  many  cubic  centimeters  of  i -normal  barium  chloride  solution 
will  be  required  to  precipitate  the  sulphate  from  one  gram  of  a  sample  con- 
taining 60%  of  Fe2(S04)3  and  14%  of  K2S04? 

5.  What  are  the  correct  conditions  for  the  precipitation  of  the  sulphate? 
Explain  in  full. 

6.  What  precautions  should  be  observed  in  the  ignition  of  the  barium 
sulphate,  and  why? 

7.  What  other  substances  may  be  determined  as  insoluble  sulphates? 
What  reagent  is  used  in  their  precipitation? 

The  Determination  of  Sulphur  in  Ores. 

1.  Assuming  complete  oxidation,  write  an  equation  to  show  the  action 
of  nitric  acid  upon  iron  pyrites,  FeS2. 

2.  If  pure  FeS2  were  decomposed  in  this  way  with  nitric  acid,  the  iron 
precipitated  with  ammonia,  and  the  filtrate  evaporated  with  hydrochloric 
acid  on  the  hot  plate  to  dryness  (to  remove  the  nitrates) ,  would  there  be  any 
danger  of  losing  a  portion  of  the  sulphur?    Explain. 

3.  Explain  the  solubility  of  lead  chloride  in  ammonium  chloride  solution. 
(Cf.  the  behavior  of  silver  chloride.) 

4.  In  the  analysis  of  an  ore  containing  lead,  how  may  we  prevent  the 
precipitation  of  a  portion  of  the  sulphur  as  lead  sulphate?    Fully  explain 
your  answer. 

5.  Outline  an  experimental  procedure  for  the  determination  of  sulphur 
in  heavy  spar,  BaS04. 


QUESTIONS  181 

The  Determination  of  Phosphoric  Anhydride  in  Phosphate  Rock. 

1.  What  are  the  chief  components  of  phosphate  rock?    What  other 
compounds  are  usually  present  ?    What  is  apatite  ? 

2.  Explain  by  means  of  the  ionic  theory  and  the  solubility  product  law 
the  fact  that  calcium  phosphate,  Ca3(P04)2,  will  dissolve  in  nitric  acid. 

3.  Why  is  it  necessary  to  remove  any  soluble  silicic  acid  that  may  be 
present  in  the  nitric  acid  solution  of  the  mineral  ?    How  is  this  done  ? 

4.  Explain  why  it  is  possible,  in  washing  the  insoluble  residue  of  silica, 
to  make  the  test  for  calcium  hi  the  wash  water  by  the  addition  of  ammonia 
alone. 

5.  Why  is  it  directed  to  neutralize  the  nitric  acid  solution  with  ammonia 
and  then  to  make  it  slightly  acid  with  nitric  acid,  before  the  addition  of  the 
molybdate  reagent  ? 

6.  Why  is  an  acid  solution  of  ammonium  nitrate  used  in  washing  the 
yellow  precipitate?     (Cf.  question  9.)    What  would  happen  if  this  am- 
monium nitrate  wash  liquid  were  alkaline  with  ammonia?    Explain. 

7.  Could  the  phosphorus  be  determined  by  igniting  and  weighing  the 
yellow  precipitate? 

8.  Why  do  we  first  precipitate  the  phosphate  with  molybdate  solution 
instead  of  precipitating  it  directly  from  the  original  solution  with  magnesia 
mixture  ? 

9.  Explain  why  a  large  excess  of  the  molybdate  reagent  is  necessary 
for  the  complete  precipitation  of  the  phosphate. 

10.  Show  by  means  of  an  equation  the  reaction  between  ammonium 
phosphomolybdate,    (NH4)3P04  .  12  MoOs,    and   ammonium  hydroxide. 
Sodium  hydroxide. 

11.  What  is  magnesia  mixture?    What  is  the  purpose  of  the  ammonium 
chloride  ?    Explain. 

12.  Why  should  the  magnesium  ammonium  phosphate  precipitation 
mixture  be  allowed  to  stand  for  some  hours  before  filtering  ? 

13.  Why  is  dilute  ammonia  used  in  washing  the  precipitate  of  mag- 
nesium ammonium  phosphate  ?    Why  not  use  water  ? 

14.  What  precautions  should  be  observed  in  the  ignition  of  this  pre- 
cipitate, and  why  ? 

15.  Name  another  element  which  may  be  determined  by  precipitation 
with  magnesia  mixture.     Can  any  elements  be  determined  in  a  similar 
manner  by  precipitation  with  sodium  or  ammonium  phosphate?    If  so, 
name  them. 

1 6.  If  it  were  desired  to  determine  the  phosphorus  in  a  sample  of  steel, 
in  which  it  is  present  as  iron  phosphide,  what  would  be  the  procedure? 
Explain. 


182     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

The  Analysis  of  Limestone. 

1.  What  is  the  principal  component  of  limestone?    What  other  com- 
pounds are  usually  present  ?    Why  is  it  important  to  analyze  a  limestone  ? 

2.  What  is  the  reason  for  the  double  evaporation  with  hydrochloric 
acid,  followed  by  continued  heating  on  the  steam  bath  ?    When  this  residue 
is  extracted  with  hydrochloric  acid,  what  is  the  insoluble  material  that  is 
left? 

3.  Why  is  the  insoluble  residue  first  washed  with  dilute  acid  rather 
than  with  water?    Would  cold  water  do  about  as  well?    Why? 

4.  What  is  the  purpose  of  adding  bromine  water  in  the  precipitation 
with  ammonium  hydroxide?    Why  must  a  large  excess  of  ammonia  be 
avoided  ? 

5.  Why  should  the  ammonium  hydroxide  precipitate  be  filtered  off 
promptly?    Of  what  does  this  precipitate  consist ?    How  is  this  precipitate 
treated,  and  why? 

6.  How  is  the  first  ammonium  oxalate  precipitate  treated,  and  why? 

7.  Explain  why  so  little  ammonium  oxalate  solution  is  added  in  the 
second  precipitation  of  the  calcium. 

8.  Explain  the  solubility  of  calcium  oxalate  in  hydrochloric  acid  and 
its  insolubility  in  acetic  acid. 

9.  What  reactions  take  place  when  calcium  oxalate,  CaC204 .  H20, 
is  ignited  ? 

10.  Why  should  the  solution  be  only  faintly  ammoniacal  for  the  pre- 
cipitation of  the  magnesium?    Why  is  more  ammonia  later  added  to  the 
precipitation  mixture?    Why  is  the  precipitation  mixture  then  allowed 
to  stand  for  several  hours  before  filtration? 

11.  How  is  the  first  precipitate  of  magnesium  ammonium  phosphate 
treated?    Explain  why  this  is  necessary.    Is  it  possible  to  accomplish  the 
same  result  in  any  other  way,  and  if  so  how  ? 

12.  Why  is  the  precipitate  washed  with  dilute  ammonia  rather  than 
with  water? 

13.  Write  an  equation  to  show  what  happens  when  MgNH4P04 .  6  H20 
is  ignited ;  when  Mg[(NH4)2P04]2  .  n  H2O  is  ignited. 

14.  What  precautions  should  be  observed  in  igniting  the  precipitate, 
and  why? 

15.  Explain  briefly  the  method  described  in  the  procedure  for  the  deter- 
mination of  carbon  dioxide. 

1 6.  What  are  the  objections  to  this  method?    Under  what  conditions 
can  the  method  be  relied  upon  to  furnish  accurate  results  ? 

17.  What  modifications  should  be  made  in  the  procedure  if  the  sample 
to  be  analyzed  is  a  baking  powder  ? 


QUESTIONS  183 

1 8.  Explain  the  operation  of  the  aspirator.    Why  is  the  carbon  dioxide 
more  readily  removed  when  air  is  bubbled  through  the  solution?     (Cf. 
Part  I,  The  Evaporation  of  Liquids.) 

19.  What  method  for  the  determination  of  carbon  dioxide  may  be  re- 
garded as  the  converse  of  the  method  described  in  the  procedure?    De- 
scribe it. 

20.  Can  you  think  of  any  other  method  for  the  exact  determination  of 
carbon  dioxide? 

The  Determination  of  Silica  in  a  Refractory  Silicate. 

1.  Why  is  it  essential  to  grind  the  whole  of  the  sample  very  fine? 

2.  Write  an  equation  to  show  the  reaction  between  orthoclase,  KAlSisOs, 
and  sodium  carbonate,  above  the  melting  point  of  the  latter.    Why  is  a 
very  large  excess  of  the  latter  used  ? 

3.  How  can  you  tell  whether  the  decomposition  is  complete  (a)  by  in- 
specting the  mixture  during  fusion;  (6)  after  the  treatment  of  the  melt 
with  dilute  acid  ? 

4.  Why  is  the  fused  mass  treated  with  a  considerable  volume  of  dilute 
acid,  rather  than  with  concentrated  acid? 

5.  How  is  the  silica  separated  from  the  residue  left  upon  evaporation 
with  dilute  hydrochloric  acid  ?    Explain. 

6.  What  is  the  purpose  of  treating  the  silica  with  hydrofluoric  and  sul- 
phuric acids  ?    Explain  in  full. 

7.  How  would  you  determine  the  percentage  of  mixed  iron  and  aluminum 
oxides  in  an  insoluble  silicate?    How  would  you  determine  the  calcium 
oxide?    The  magnesium  oxide? 

8.  How  would  you  determine  the  silica  in  a  silicate  mineral  which  may 
be  readily  decomposed  by  hydrochloric  acid  ? 

The  Determination  of  Potash. 

1.  Is  this  a  precipitation  method ?    If  not,  to  what  class  does  it  belong? 

2.  Why  is  it  necessary  to  remove  hydrochloric  and  sulphuric  acids  before 
the  extraction  of  the  soluble  perchlorates  ?    How  are  these  acids  removed  ? 
Explain. 

3.  What  weight  of  HC104  will  be  required  to  convert  one  gram  of  NaCl 
into  NaC104?    One  gram  of  KC1  into  KC1O4?    What  connection  exists 
between  these  questions  and  the  procedure?     (Cf.  Part  IV.,  Problem  30.) 

4.  Explain  the  fact  that  phosphates,  even  if  insoluble  in  alcohol,  need 
not  be  removed  before  making  the  extraction.    What  precaution  should  be 
observed  when  phosphates  are  present,  and  why? 

5.  Do  any  of  the  following  salts  interfere  with  the  method ;  and,  if  so, 
how  is  the  difficulty  overcome?    NH4C104,  Ba(C104)2,  Mg(C104)2. 


184     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

6.  Why  is  a  small  quantity  of  HC1O4  added  to  the  alcohol  to  be  used  in 
the  extraction?    If  50  cc.  of  this  liquid,  capable  of  dissolving  2  mg.  of  pure 
KC104,  are  used  in  the  extraction,  why  is  not  the  quantity  of  KC104  found 
in  the  residue  2  mg.  less  than  the  true  value  ? 

7.  How  may  we  determine  the  point  at  which  the  extraction  is  complete  ? 
Explain. 

The  Electrolytic  Determination  of  Copper. 

1.  Define:  ampere;  volt;  ohm.    State  Ohm's  law.    State  Faraday's 
laws. 

2.  Show  by  means  of  a  diagram  how  the  apparatus  is  assembled  for 
an  electrolytic  determination. 

3.  Explain  the  electrolytic  reduction  of  nitric  acid  to  ammonia. 

4.  Explain  the  deposition  of  lead  peroxide  upon  the  anode,  in  the 
electrolysis  of  a  solution  containing  lead  nitrate  and  nitric  acid. 

5.  What  is  meant  by  the  discharge  potential  of  an  ion?    The  decom- 
position voltage  of  a  salt  ?    What  is  polarization  ?    Explain  in  each  case. 

6.  Explain  why  it  is  possible  to  separate  copper  from  nickel  by  elec- 
trolysis?    Can  nickel  be  separated  from  cobalt  in  this  way?    Why? 

7.  What  are  the  advantages  of  mechanical  stirring  during  electrolysis? 

8.  What  is  meant  by  current  density?    What  is  the  unit  of  current 
density?    Why  is  current  density  a  factor  of  the  greatest  importance  in 
electro-analysis  ? 

9.  What  factors  favor  the  formation  of  a  satisfactory  deposit?    What 
factors  interfere  with  it  ? 

10.  Outline  a  method  for  the  preparation  of  a  solution  suitable  for 
electrolysis  when  the  sample  to  be  analyzed  is  a  copper  ore  containing 
arsenic.    Why  should  the  arsenic  be  removed  ? 

11.  Discuss  the  materials  from  which  electrodes  may  be  prepared,  and 
the  form  of  the  electrodes  and  electrolytic  vessels. 

12.  Name  the  factors  in  the  electrolytic  work  which  should  receive 
especial  attention  in  an  endeavor  to  make  accurate  and  reliable  copper 
determinations  as  rapidly  as  possible.    Explain  in  the  case  of  each  factor 
mentioned. 

Volumetric  Analysis :  Fundamental  Principles. 

1.  Specify  the  chief  uses  of  measuring  flasks;  of  transfer  pipettes;  of 
burettes. 

2.  Describe  the  preparation  of  sulphuric  acid-dichromate  cleaning  solu- 
tion.   How  is  it  used? 

3.  In  a  titration  with  a  solution  correctly  made  up  to  tenth-normal  con- 
centration at  20°,  a  burette  correctly  graduated  for  use  at  20°  is  used  at  an 


QUESTIONS  185 

actual  temperature  of  27.5°,  and  the  indicated  volume  of  solution  withdrawn 
is  27.68  cc. ;  to  how  many  cubic  centimeters  of  tenth-normal  solution  does 
this  liquid  correspond  ?  (See  Part  I.) 

4.  Define  the  term  "liter." 

5.  Describe  a  method  for  the  calibration  of  a  loo-cc.  measuring  flask. 
Illustrate  your  description.     (See  Part  I.) 

6.  What  is  a  standard  solution?    A  normal  solution?    Define  and 
illustrate  the  term  "  normality  factor." 

7.  Characterize  in  general  the  reactions  which  are  suitable  as  the  basis 
for  volumetric  processes. 

8.  What  is  an  indicator?    Illustrate  your  answer. 

9.  Discuss  the  advantages  of  the  volumetric  system. 

Neutralization  Methods :  The  Standardization  of  Acids  and  Alkalies. 

1.  Define  in  terms  of  the  theory  of  ionization:   (a)  a  neutral  solution; 
(6)  an  acid  solution ;  (c}  an  alkaline  solution. 

2.  Will  the  solution  resulting  from  mixing  equal  volumes  of  one-tenth 
normal  aqueous  solutions  of  the  following  substances  be  acid,  alkaline,  or 
neutral:   (a)    hydrogen    chloride    and    sodium    hydroxide;  (6)    hydrogen 
chloride  and  ammonium  hydroxide ;  (c)  acetic  acid  and  sodium  hydroxide  ? 

3.  Give  a  full  explanation  of  case  (c)  above,  writing  all  equations  and 
equilibria. 

4.  What  indicator  should  be  used  in  each  case  in  titrations  involving 
the  combinations  indicated  in  question  2  ?    Explain  fully.     (See  the  section 
on  Indicators  for  Use  in  Alkalimetry  and  Acidimetry.) 

5.  Outline  procedures  for  the  preparation  of  approximately  half -normal 
solutions  of  hydrochloric  acid  and  sodium  hydroxide. 

6.  Describe  a  method  for  obtaining  the  ratio  between  the  solutions 
referred  to  in  the  preceding  question.    If  20.00  cc.  of  an  acid  solution  re- 
quire 21.46  cc.  of  0.4693 -normal  alkali  for  neutralization,  what  is  the  nor- 
mality factor  of  the  acid  ? 

7.  Describe  the  sodium  carbonate  method  for  the  standardization  of  a 
solution  of  hydrochloric  acid.    If  0.5383  g.  of  Na2C03  require  20.15  cc.  of 
the  acid  for  neutralization,  what  is  the  normality  factor  of  the  acid? 

8.  If  10.00  cc.  of  a  solution  of  hydrogen  chloride  yield  0.7421  g.  of  silver 
chloride,  what  is  the  normality  factor  of  the  acid  ? 

9.  Describe  the  standardization  of  sodium  hydroxide  solution  by  means 
of  potassium  bitartrate.     If  1.179  g-  of  tne  latter  require  13.35  cc-  of  the 
alkali  for  neutralization,  what  is  the  normality  factor  of  the  solution? 

10.   If  20.00  cc.  of  an  acid  are  equivalent  to  21.20  cc.  of  an  alkali,  and  if 
40.00  cc.  of  the  acid  are  added  to  0.6000  g.  of  Na2C03  and  the  resulting  solu- 


1 86     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

tion  requires  3.00  cc.  of  the  alkali  for  neutralization,  what  are  the  normality 
factors  of  the  two  solutions  ? 

ii.  Describe  a  method  for  the  preparation  of  a  solution  of  hydrochloric 
acid  of  exactly  one-half  normal  concentration. 

The  Total  Alkalinity  of  Soda  Ash. 

1.  Discuss  the  composition  of  soda  ash,  and  account  for  the  impurities 
it  is  likely  to  contain. 

2.  Which  of  the  impurities  will  contribute  to  the  total  alkaline  strength 
of  the  soda  ash? 

3.  Could  this  determination  be  made  with  phenolphthalein  as  the  indi- 
cator, and  if  so  how  ? 

4.  Give  two  volumetric  methods  for  the  determination  of  Na2C03  and 
of  NaOH  in  mixtures  of  the  two  substances. 

5.  Give  a  method  for  the  determination  of  Na2C03  and  of  NaHC03  in 
mixtures  of  the  two  salts. 

The  Neutralization  Value  of  an  Acid. 

1.  In  titrating  weak  acids,  why  is  it  usually  necessary  to  work  at  the 
boiling  temperature? 

2.  At  what  point  does  an  ice-cold,  dilute  solution  of  sodium  hydroxide 
containing  phenolphthalein  lose  its  red  color  upon  being  treated  with  carbon 
dioxide?    Upon  boiling  this  decolorized  solution,   the  color  reappears; 
explain  fully  the  mechanism  by  which  the  alkalinity  of  the  solution  increases 
on  boiling. 

3.  Describe  a  method  for  the  preparation  of  a  carbonate-free  solution 
of  sodium  hydroxide.    How  should  such  a  solution  be  preserved?    Why  is 
its  preparation  sometimes  worth  while  ? 

4.  Is  it  possible  to  accurately  titrate  sulphurous  acid  with  a  standard 
solution  of  ammonia?    If  not,  why?    And  if  so,  what  indicator  should  be 
used? 

The  Determination  of  Protein  Nitrogen  by  the  Kjeldahl  Method. 

1.  What  is  the  purpose  in  the  digestion  of  (a)  the  concentrated  sul- 
phuric acid?     (6)  The  copper  sulphate?     (c)  The  potassium  sulphate? 

2.  What  chemical  change  does  sulphuric  acid  undergo  during  the  di- 
gestion ?    The  organic  matter  ? 

3.  Why  is  a  long-necked  flask  used  ? 

4.  How  would  the  procedure  of  digestion  be  modified  if  nitrates  were 
present  ?    Explain  fully. 

5.  How  would  the  procedure  be  modified  if  mercury  were  added  instead 
of  copper  sulphate  ?    Explain  in  full. 


QUESTIONS  187 

6.  In  what  form  does  the  nitrogen  exist  after  the  completion  of  the 
digestion?    After  making  the  solution  alkaline  with  sodium  hydroxide? 
In  what  form  does  it  distill  over  ? 

7.  How  is  the  loss  of  ammonia  prevented  upon  the  addition  of  an  excess 
of  sodium  hydroxide  ? 

8.  What  is  the  purpose  of  the  zinc?    What  is  the  action  of  sodium 
hydroxide  solution  upon  zinc  ? 

9.  The  ammonia  from  one  gram  of  a  fertilizer  is  distilled  into  20.00  cc. 
of  0.5000  ^V  acid,  and  6.00  cc.  of  0.4800  N  alkali  are  required  to  neutralize 
the  excess  of  acid ;  calculate  the  percentage  of  nitrogen  in  the  sample. 

10.  If  0.20  cc.  of  25.00%  NaOH  (sp.  gr.,  1.25)  had  been  carried  over  by 
bumping  or  foaming,  what     ould  have  been  the  apparent  percentage  of 
nitrogen  in  the  above  case? 

11.  If  you  had  to  determine  the  percentage  of  NH3  and  of  HCzH-sOz  in 
crude  ammonium  acetate,  how  would  you  proceed? 

Bichromate  Methods :  The  Titration  of  Iron. 

1.  Write  the  equation  for  the  oxidation  of  ferrous  chloride  with  potas- 
sium dichromate  in  the  presence  of  hydrochloric  acid. 

2.  What  weight  of  K2Cr2O7  is  required  for  one  liter  of  a  tenth-normal 
solution,  to  be  used  as  an  oxidizing  agent  ? 

3.  Outline  the  procedure  for  the  standardization  of  dichromate  solu- 
tion by  means  of  iron  wire.    Why  is  it  well  also  to  have  a  standard  solution 
of  ferrous  ammonium  sulphate,  and  how  is  this  solution  standardized? 

4.  What  is  the  maximum  weight  of  pure  iron  wire  which  can  be  taken 
for  reaction  with  tenth-normal  dichromate  without  having  to  refill  a  3o-cc. 
burette? 

5.  Name  four  reagents  which  can  be  used  to  reduce  ferric  salts  to  ferrous, 
in  the  presence  of  hydrochloric  acid,  and  write  the  equation  in  each  case. 

6.  Why  is  it  necessary  after  reduction  with  stannous  chloride  to  add 
mercuric  chloride  to  the  solution?    Why  must  the  stannous  chloride  be 
present  only  in  very  slight  excess  (equations)  ?    Why  should  hydrochloric 
acid  be  present  during  the  titration  ? 

7.  What  indicator  is  used  in  the  titration?    What  action  has  it  upon 
ferric  salts?     Upon  ferrous  salts? 

8.  Why  is  the  indicator  used  outside  of  the  solution? 

9.  What  products  are  formed  when  chromite,  Fe(CrO2)2,  is  fused  with 
sodium  peroxide  ?    What  happens  upon  treating  the  fused  mass  with  water  ? 
Why  is  the  aqueous  solution  boiled?    What  then  happens  when  the  filtrate 
is  acidified  with  sulphuric  acid?    How  may  the  chromium  in  the  resulting 
solution  be  determined  volumetrically  ? 


i88     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Permanganate  Methods :  The  Titration  of  Iron  and  of  Oxalic  Acid. 

1 .  How  many  grams  of  KMnC>4  are  required  for  one  liter  of  tenth-normal 
solution,  to  be  used  as  an  oxidizing  agent  in  acid  solutions  ?     To  be  used  in 
neutral  solutions  for  the  oxidation  of  manganese? 

2.  Why  should  a  permanganate  solution  be  allowed  to  stand  for  several 
days,  and  then  be  filtered  through  asbestos,  before  it  is   standardized? 
Why  should  it  not  be  placed  in  burettes  having  rubber  outlet  tubes? 

3.  Name  at  least  four  substances  which  can  be  used  to  standardize 
permanganate  solutions.    Write  an  equation  in  each  case. 

4.  What  is  the  maximum  weight  of  Na2C204  which  can  be  titrated 
with  tenth-normal  permanganate  solution  without  having  to  refill  a  30-cc. 
burette? 

5.  In  the  titration  of  oxalic  acid,  why  is  it  that  the  oxidation  proceeds 
so  much  more  slowly  at  first  than  later  on?    Explain  in  full. 

6.  Name  ten  substances  which  can  be  quantitatively  determined  by 
means  of  potassium  permanganate. 

7.  What  effect  does  potassium  permanganate  have  upon  hydrochloric 
acid  in  the  presence  of  ferrous  salts,  even  in  very  dilute  solution?    How 
may  this  action  be  prevented  ? 

8.  What  are  the  components  of  the  Zimmermann-Reinhardt  solution? 
Explain  the  purpose  of  each. 

9.  Under  what  conditions  can  ferrous  iron  be  accurately  determined 
with  potassium  permanganate  without  the  use  of  the  Zimmermann-Rein- 
hardt solution? 

10.  What  is  the  maximum  weight  of  a  sample  of  ore  containing  40.00% 
of  iron  which  can  be  taken  for  titration  with  tenth-normal  oxidizing  agent 
without  having  to  refill  a  3o-cc.  burette  ? 

11.  Discuss  the  preparation  of  a  solution  for  analysis  from  a  refractory 
iron  ore. 

12.  Describe  a  method  for  the  determination  of  calcium  by  means  of 
potassium  permanganate. 

13.  Outline  a  method  for  the  determination  of  the  oxidizing  power  of 
pyrolusite  by  means  of  oxalic  acid  and  potassium  permanganate. 

14.  How  can  the  determination  referred  to  in  the  preceding  question  be 
made  by  a  gravimetric  process  ? 

The  Determination  of  Phosphorus  in  Steel. 

1.  Assuming  the  presence  of  the  phosphorus  as  FesP2,  show  by  means 
of  an  equation  the  action  of  nitric  acid  upon  this  compound.    Why  is  the 
nitric  acid  solution  heated  with  potassium  permanganate? 

2.  In  order  to  cause  the  higher  oxides  of  manganese,  such  as  MnOgj  to 


QUESTIONS  189 

go  into  solution  in  nitric  acid,  what  kind  of  a  reagent  should  be  added? 
Illustrate  and  explain. 

3.  What  is  the  purpose  of  precipitating  the  phosphorus  as  ammonium 
phosphomolybdate  ? 

4.  Why  is  the  yellow  precipitate  dissolved  in  ammonia  and  the  solution 
acidified  with  sulphuric  acid,  rather  than  to  dissolve  it  directly  in  sulphuric 
acid?    Why  is  Mo03  not  precipitated  when  the  ammoniacal  solution  is 
a  cidified  with  sulphuric  acid  ?    Would  hydrochloric  or  nitric  acid  do  as  well 
here,  and  why? 

5.  Describe  the  construction  and  use  of  the  Jones  reductor.    Can  it  be 
used  for  the  reduction  of  substances  other  than  molybdenum? 

6.  Why  is  it  best  to  receive  the  reduced  molybdenum  solution  below 
the  surface  of  a  solution  containing  ferric  alum?    What  is  the  purpose  of 
the  phosphoric  acid  in  this  solution? 

7.  How  should  the  permanganate  solution  be  standardized  hi  order  to 
obtain  the  most  reliable  results? 

8.  Would  you  recommend  the  determination  of  phosphoric  anhydride 
in  apatite  by  this  method  ?    Why  ?    Is  there  a  suitable  volumetric  method  ? 
If  so,  describe  it. 

The  Determination  of  Manganese  in  an  Ore. 

1.  What  is  the  Guyard  reaction  ?    What  role  does  it  play  in  the  titration 
of  oxalic  acid  or  of  iron  with  potassium  permanganate? 

2.  Discuss  the  preparation  of  the  solution  for  analysis  from  a  refractory 
ore  containing  manganese. 

3.  What  happens  when  zinc  oxide  is  added  to  the  acid  solution  of  the 
ore  (equations)  ? 

4.  Explain  fully  why  a  zinc  salt  should  be  present  in  the  solution  during 
the  titration. 

5.  Explain  why  the  presence  of  chlorides  does  not  interfere  with  the 
accuracy  of  this  titration. 

6.  If  the  permanganate  solution  used  in  this  titration  is  o.iooo-N  for 
use  with  iron  or  oxalic  acid,  what  is  its  normality  factor  for  this  reaction? 

7.  How  is  it  best  to  standardize  the  permanganate  solution  used  hi  this 
determination  ? 

lodometric  Methods :  The  Preparation  and  Standardization  of  Iodine 
and  Thiosulphate  Solutions. 

1.  How  many  grams  of  iodine  are  required  for  one  liter  of  the  tenth- 
normal  solution?    Of  sodium  thiosulphate,  Na2S203 .  5  H20? 

2.  How  is  the  iodine  solution  made,  and  what  is  the  purpose  of  the 
potassium  iodide  ?    Show  what  equilibria  exist  in  the  iodine  solution.    Why 


IQO     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

should  the  water  be  freshly  boiled  and  allowed  to  cool  out  of  contact  with 
the  air,  in  the  preparation  of  the  thiosulphate  solution  ? 

3.  Write  the  reaction  between  iodine  and  sodium  thiosulphate.    How 
do  chlorine  and  bromine  differ  from  iodine  in  their  behavior  towards  sodium 
thiosulphate?    Explain  why  this  is  so. 

4.  What  is  the  effect  of  free  carbonic  acid  upon  sodium  thiosulphate 
solution?    Does  the  decomposition  cease  as  soon  as  all  of  the  carbonic 
acid  has  reacted?    Write  equations  to  illustrate  your  answers. 

5.  Does  the  solution  resulting  from  the  partial  decomposition  of  the 
thiosulphate  have  a  greater  or  lower  reducing  value  than  the  original  solu- 
tion?   Explain  why. 

6.  Give  equations  to  show  two  ways  in  which  iodine  may  act  as  an 
oxidizing  agent. 

7.  What  is  the  maximum  weight  of  As20a  which  can  be  taken  for  reaction 
with  tenth-normal  iodine  solution  without  having  to  refill  a  30-cc.  burette? 

8.  Can  the  standardization  of  iodine  against  arsenious  oxide  be  per- 
formed in  a  strongly  alkaline  solution  ?    Can  it  be  done  in  an  acid  solution  ? 
Give  reasons  for  your  answers. 

9.  What  is  the  purpose  of  the  sodium  bicarbonate  ?    Is  the  bicarbonate 
solution  acid,  alkaline,  or  neutral?    Explain  your  answer. 

10.  Discuss  the  determination  of  the  end-point.    Explain  why  the  indi- 
cator is  added  hi  such  large  quantity. 

11.  Discuss  the  use  of  iodine  solutions  in  the  presence  of  sulphuric  acid. 
In  the  presence  of  ammonium  salts. 

The  Determination  of  Antimony  in  Stibnite. 

1.  Write  the  reaction  between  pure  stibnite  and  hydrochloric  acid. 

2.  Why  must  the  hydrochloric  acid  solution  be  heated  on  the  steam 
bath  ?    Why  must  it  not  be  boiled  until  after  dilution  ? 

3.  What  is  the  purpose  of  adding  tartaric  acid  to  the  solution  ?    Explain. 

4.  Explain  why  the  solution  may  possibly  turn  red  during  gradual 
dilution.    What  is  the  correct  procedure  in  such  a  case? 

5.  If  a  white  precipitate  forms  upon  dilution,  what  error  has  been  made  ? 
What  is  the  white  precipitate,  and  what  should  be  done  with  the  mixture? 

6.  What  is  the  purpose  of  almost  neutralizing  the  solution  with  sodium 
hydroxide,  and  how  is  this  accomplished?    Why  is  sodium  bicarbonate 
then  added  in  large  excess  ? 

7.  What  elements  would,  if  present,  interfere  with  this  determination, 
and  why  ? 

The  Determination  of  Lead  in  an  Ore. 

i.  After  the  decomposition  of  the  ore  and  the  addition  of  sulphuric  acid, 
why  is  it  necessary  to  evaporate  to  white  fumes? 


QUESTIONS  191 

2.  Explain  the  solubility  of  lead  sulphate  in  ammonium  acetate  solution. 

3.  Explain  why  it  is  possible  to  quantitatively  precipitate  the  lead  from 
the  ammonium  acetate  solution  by  means  of  an  excess  of  potassium  dichro- 
mate. 

4.  Explain  the  solubility  of  lead  chromate  in  the  acidified  solution  of 
sodium  chloride. 

5.  Write  an  equation  to  show  the  reaction  of  the  acid  chromate  solution 
with  potassium  iodide. 

6.  Why  should  the  thiosulphate  solution  used  in  this  determination  be 
standardized  under  identical  conditions  against  test  lead  ? 

The  Determination  of  Copper  in  an  Ore. 

1.  In  this  determination,  why  is  the  copper  separated  from  the  other 
metals  present  in  the  ore?    Explain  how  iron,  arsenic,  or  antimony  would 
interfere  with  the  accuracy  of  the  titration. 

2.  What  would  you  expect  the  composition  of  the  precipitate  to  be 
which  is  formed  upon  the  addition  of  sodium  thiosulphate  to  a  solution  of 
copper  sulphate?    How  does  it  happen,  then,  that  we  obtain  cuprous 
sulphide  ? 

3.  Can  any  other  metals  be  precipitated  from  their  salt  solutions  by 
means  of  sodium  thiosulphate?     (Try,  for  example,  silver  nitrate  and 
sodium  thiosulphate,  in  the  cold,  and  explain  what  takes  place.) 

4.  What  is  the  object  of  igniting  the  precipitate  of  cuprous  sulphide? 
What  becomes  of  any  antimony  which  is  present  ? 

5.  Why  is  it  so  important  to  standardize  the  thiosulphate  solution 
against  pure  metallic  copper? 

6.  Why  must  the  nitrous  fumes  be  completely  expelled  before  the  ad- 
dition of  the  potassium  iodide  ? 

7.  Why  is  it  preferable  to  titrate  the  free  iodine  in  the  presence  of  acetic 
acid,  rather  than  in  the  presence  of  sulphuric  acid?    Explain  fully. 

8.  Write  an  equation  to  show  the  action  of  nitric  acid  upon  cuprous 
iodide. 

Precipitation  Methods :  The  Determination  of  Chlorine. 

1.  Briefly  outline  the  procedure  for  the  standardization  of  the  silver 
nitrate  and  potassium  thiocyanate  solutions  against  pure  sodium  chloride. 
Could  these  solutions  be  standardized  against  pure  metallic  silver,  and  if 
so  how  ? 

2.  How  may  pure  sodium  chloride  be  prepared  from  the  commercial  salt  ? 

3.  What  indicator  is  used  in  connection  with  thiocyanate  solutions? 
Why  must  nitric  acid  be  present?    Explain  fully  why  the  indicator  should 
be  added  in  such  large  quantity. 


INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

4.  At  18-20°,  the  solubility  product  of  silver  chloride  is  about  o.6Xio~10 
and  that  of  silver  thiocyanate  is  about  o.6Xio~12;  what  is  the  relative  con- 
centration of  the  chloride  and  thiocyanate  ions  in  a  solution  which  is  satu- 
rated with  both  salts?    Assume  the  equal  (practically  complete)  ioniza- 
tion  of  both  salts. 

5.  Why  is  it  necessary  to  filter  off  the  silver  chloride  before  making  the 
titration  with  the  thiocyanate  solution?    Base  your  explanation  upon  the 
data  given  in  the  preceding  question. 

6.  In  general,  what  anions  may  be  determined  by  this  method  without 
first  filtering  off  the  silver  salt? 

7.  How  may  the  halogens  in  alkali  chlorates,  bromates,  and  iodates  be 
determined  by  this  method  ? 

8.  Outline  a  procedure  for  the  determination  by  this  method  of  (a)  the 
chlorine  in  horn  silver,  AgCl ;  (6)  the  silver  in  the  same  mineral. 


APPENDIX 

THE  PREPARATION  OF   THE  REAGENTS 

MANY  of  the  reagents  used  in  quantitative  analysis  are  pre- 
pared for  one  specific  purpose  only,  and  directions  for  the  prepa- 
ration of  such  reagents  will  be  found  in  the  treatment- of  the 
individual  determinations.  Certain  reagents,  however,  are 
used  at  approximately  fixed  concentrations  in  a  variety  of 
processes,  and  it  is  especially  these  which  are  included  in  this 
section. 

There  are  many  advantages  in  making  the  concentrations  of 
the  reagents  used  in  analytical  work  follow  a  definite  system. 
The  most  convenient  system  is  to  use  multiples  or  submulti- 
ples  of  the  equivalent  weight  employed  in  volumetric  analysis, 
though  of  course  the  concentration  of  the  solution  need  not 
be  known  with  such  exactitude  as  in  volumetric  work. 

With  this  system,  equal  volumes  of  the  solutions  bear  fixed 
relations  to  one  another,  it  is  easy  to  calculate  the  volume  of  a 
reagent  which  is  required  for  a  specific  purpose,  and  the  addi- 
tion of  an  unnecessary  excess  may  readily  be  avoided.  This 
means  a  saving  in  time,  labor,  and  material;  and  it  leads  to 
more  accurate  and  reliable  work.  Thus,  if  it  is  directed  to  fuse 
one  gram  of  a  silicate,  say  KAlSi308,  with  7.5  g.  of  Na2C03 
and  to  take  up  the  cooled  mass  with  water  and  an  excess  of 
hydrochloric  acid,  we  know  that  24  cc.  of  6-normal  acid  will 
suffice  to  neutralize  the  mixture  (since  7.5  g.  of  Na2C03  corre- 
spond to  about  one  seventh  of  an  equivalent  of  this  salt),  and 
that  30  cc.  in  all  will  furnish  a  sufficient  excess.  If  the  mixture 
is  evaporated  to  dryness  in  the  regular  manner  and  taken  up 
with  2  cc.  of  i2-normal  hydrochloric  acid  and  a  little  water,  it 
o  193 


194     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

will  be  obvious  that  4  cc.  of  6-normal  ammonium  hydroxide  will 
suffice  to  neutralize  the  acid ;  and  that  8  cc.  in  all  will  neutralize 
the  acid,  precipitate  the  aluminum,  and  in  addition  furnish  a 
sufficient  excess. 

Measuring  cylinders  and  measuring  pipettes  are  useful  for 
delivering  specific  volumes  of  such  reagents. 

ACIDS 

Acetic,  6-normal :  Mix  350  cc.  of  glacial  acetic  acid  with 
650  cc.  of  water. 

Hydrochloric,  i2-normal:  Use  the  C.  P.  acid  of  commerce 
of  sp.  gr.,  1.19. 

Hydrochloric,  6-normal :  Mix  12 -normal  acid  with  an  equal  vol- 
ume of  water.  The  specific  gravity  of  this  acid  is  about  i.io. 

Nitric,  1 6-normal:  Use  the  C.  P.  acid  of  commerce  of  sp.  gr., 
1.42. 

Nitric,  6-normal :  Mix  380  cc.  of  the  1 6-normal  acid  with  650 
cc.  of  water.  The  specific  gravity  of  this  acid  is  about  1.195. 

Sulphuric,  3 6-normal :  Use  the  C.  P.  acid  of  commerce  of  sp. 
gr.,  1.84. 

Sulphuric,  6-normal:  Pour  200  cc.  of  the  3 6-normal  acid 
into  1045  cc.  of  water.  The  specific  gravity  of  this  acid  is  1.18. 

BASES 

Ammonium  hydroxide,  i5-normal:  Use  the  C.  P.  ammonia 
water  of  commerce  of  sp.  gr.,  0.90. 

Ammonium  hydroxide,  6-normal:  Mix  400  cc.  of  the  15- 
normal  solution  with  600  cc.  of  water.  The  specific  gravity 
of  this  solution  is  about  0.958. 

Sodium  hydroxide,  6-normal:  Dissolve  250  g.  of  stick  sodium 
hydroxide  in  water  and  dilute  to  one  liter. 

SALTS 

Ammonium  carbonate:  Dissolve  250  g.  of  freshly  powdered 
ammonium  carbonate  in  one  liter  of  6-normal  ammonium  hy- 
droxide, and  filter  if  there  is  a  residue. 


APPENDIX  195 

Ammonium  mandate?  Dissolve  100  g.  of  MoOs  in  80  cc. 
of  ammonia  (sp.  gr.,  0.90)  with  the  addition  of  400  cc.  of  water ; 
with  cooling  and  constant  stirring,  allow  the  clear  solution  to 
run  slowly  into  a  mixture  of  400  cc.  of  nitric  acid  (sp.  gr., 
1.42)  with  600  cc.  of  water,  add  0.05  g.  of  rmcrocosmic  salt, 
NaNH4HP04 . 4  H2O,  and  keep  the  mixture  in  a  warm  place 
for  several  days,  or  until  a  portion  heated  to  40°  deposits  no  yellow 
precipitate.  Decant  from  any  sediment,  and  preserve  in  glass- 
stoppered  bottles.  This  solution  contains  68  g.  of  Mo03  per  liter. 

Ammonium  oxalate,  o.5-normal:  Dissolve  35.5  g.  of 
(NH4)2C2O4 .  H20  in  1000  cc.  of  water. 

Barium  chloride,  i -normal:  Dissolve  122  g.  of  BaCl2 .  2  H20 
in  1000  cc.  of  water. 

Magnesia  mixture,  o.5-normal  as  a  precipitant  for  phosphoric 
or  arsenic  acid:  Dissolve  51  g.  of  MgCl2 .  6  H20  and  130  g. 
of  NH4C1  in  water,  add  121  cc.  of  ammonia  (sp.  gr.,  0.90),  and 
dilute  to  one  liter. 

Mercuric  chloride,  o.2-normal  for  oxidizing  stannous  chloride : 
Dissolve  54  g.  of  HgCl2  in  loco  cc.  of  hot  water. 

Silver  nitrate,  o.2-normal :  Dissolve  34  g.  of  AgN03  in  1000  cc. 
of  water. 

Sodium  phosphate,  o.5-normal  as  a  precipitant  for  mag- 
nesium: Dissolve  90  g.  of  Na2HP04 .  12  H20  (or  52  g.  of 
NaNH4HP04 .  4  H20)  in  1000  cc.  of  water. 

Stannous  chloride,  i -normal  as  a  reducing  agent:  Dissolve 
113  g.  of  SnCl2 .  2  H20  in  150  cc.  of  i2-normal  hydrochloric 
acid,  with  the  gradual  addition  of  water,  finally  diluting  to  one 
liter.  Keep  in  bottles  containing  granulated  tin. 

1  Recovery  of  the  Molybdk  Acid.  To  the  liquid  molybdate  residues,  acidified 
if  necessary  with  nitric  acid,  add  sodium  phosphate  solution  in  excess.  Collect 
the  yellow  precipitate,  wash  it  with  water  containing  sodium  sulphate,  and  then 
dry  it  in  the  air.  Dissolve  i  Kg.  of  the  dried  precipitate  in  ammonia,  add  a  strong 
solution  of  60  g.  of  NH4C1  and  120  g.  of  MgClz .  6  H2O  in  water,  allow  to  stand  for 
6  hours,  and  filter  off  the  precipitate.  To  the  filtrate,  decolorized  if  necessary  with 
a  little  H2O2,  add  HC1  just  to  acid  reaction,  to  precipitate  the  MoO3.  Collect 
this  precipitate,  wash  with  water,  and  dry  at  110°. 


196     INTRODUCTORY  COURSE  IN  QUANTITATIVE  ANALYSIS 

Sulphuric  Acid-Dichr ornate  Cleaning  Solution 

With  stirring,  cautiously  pour  200  cc.  of  sulphuric  acid  (sp. 
gr.,  1.84)  into  150  cc.  of  cold  water,  and  saturate  the  hot  solu- 
tion, without  further  heating,  with  powdered  sodium  (or  potas- 
sium) dichromate. 

When  cleaning  measuring  vessels  with  this  liquid,  they  should 
be  filled  with  the  cold  solution  and  allowed  to  stand  overnight, 
or  longer. 

Analytical  Samples  for  the  Use  of  Students 

The  analyzed  samples  indicated  in  the  text  for  the  use  of 
beginners  in  quantitative  analysis  may  in  some  cases  be  ob- 
tained in  the  market.  Otherwise,  they  may  be  prepared  by 
mixing  together  the  component  materials  in  the  proportions 
decided  upon.  This  mixing  is  best  accomplished  by  long  con- 
tinued grinding  in  a  ball  mill,  the  material  being  finally  passed 
through  a  fine-meshed  sieve,  and  bottled.  These  samples  should 
be  carefully  analyzed  by  members  of  the  quantitative  staff,  so 
that  the  student's  work  may  be  judged  according  to  its  accuracy. 
Most  of  the  mixtures  indicated  can  be  kept  from  year  to  year 
without  change. 

It  is  desirable  to  have  in  each  case  a  continuous  series  of 
at  least  ten  samples,  varying  in  content  from  sample  to  sample 
by  about  0.4-0.5%. 


APPENDIX 


197 


APPARATUS  IN  THE  STUDENT'S  DESK1 
QUANTITATIVE  CHEMICAL  LABORATORIES 


Above  in  the  drawers 
i  Brush,  camel's  hair, 
i  Burette,  g.  s.,  30  cc. 

1  Burette,  30  cc.,  for  pinchcock. 
4  Crucibles,  porcelain,  o. 

2  Crucibles,  porcelain,  Gooch,  extra 

disc. 
2  Cylinders,  graduated,  50  cc.  and 

10  cc. 
i  File. 

1  Forceps,  steel. 

2  Funnels,    diam.    25    mm.,    stem 

40  mm. 

2  Glasses,  watch,  140  mm. 
2  Glasses,  watch,  70  mm. 
2  Glasses,  watch,  50  mm. 
i  Vial  litmus  paper,  blue. 

1  Vial  litmus  paper,  red. 

2  Boxes  matches,  safety, 
i  Pinchcock. 

i  Pipette,  25  cc. 

1  Pipette,  10  cc. 

2  Policemen,  rubber  tip. 

3  Rods,  glass,  200  mm. 
6  Test  tubes. 

i  Thermometer,  100°  C. 

i  Tongs,  brass,  nickel  plated. 

1  Tube,  connecting,  3-way. 

2  Tubes,  rubber,  for  Gooch  crucibles. 

3  Tubes,  weighing,  with  corks. 

i    Tube,   rubber,   pressure,   length 

300  mm. 
i    Tube,    rubber,     small,     length 

300  mm. 
Tubing,  soft  glass,  900  mm. 

1  The  articles  listed  above  represent  the  apparatus  with  which  it  is  desirable 
to  provide  each  student  at  the  outset ;  the  list  can  of  course  be  modified  in  many 
particulars  without  jeopardizing  the  success  of  the  work.  Any  additional  appa- 
ratus which  may  be  required  can  be  obtained  as  needed  from  the  store  room. 


Below  in  the  cupboard 
12  Beakers,  2  nests,  1-6,  with  win- 
dow pane. 

2  Bottles,  g.  s.,  2500  cc. 
i  Bottle,  g.  s.,  250  cc.,  for  cleaning 

solution, 
i  Bottle,  g.  s.,  125  cc.,  for  silver 

nitrate, 
i  Bottle,  weighing. 

1  Burette  holder,  Lincoln's. 

2  Burners,  adjustable. 

2  Burner  tubes,  rubber. 

2  Casseroles,  porcelain,  500  cc. 

1  Desiccator  for  4  crucibles. 

2  Flasks,  Erlenmeyer,  500  cc. 
2  Flasks,  Erlenmeyer,  250  cc. 
2  Flasks,  Erlenmeyer,  150  cc. 
2  Flasks,  filter,  500  cc. 

2  Flasks,  Florence,  500  cc. 
2  Flasks,  Florence,  250  cc. 
2  Flasks,  Florence,  50  cc.,  for  in- 
dicators, 
i  Flask,  volumetric,  1000  cc. 

1  Flask,  volumetric,  500  cc. 

2  Flasks,  volumetric,  250  cc. 

4  Funnels,   diam.   70   mm.,   stem 

200  mm. 

2  Funnels,  for  Gooch  crucibles, 
i  Sponge. 

1  Stand,  filter,  wooden. 

2  Stands,  iron,  i  ring  each. 

2  Triangles,  pipe  stem,  new  form. 
2  Tripods,  iron. 
2  Wire  gauzes. 


198 


Table  A — Four  Place  Logarithms 


N 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

10 

0000 

0043 

0086 

0128 

0170 

0212 

0253 

0294 

0334 

0374 

4   812 

17  21  25 

29  33  37 

11 
12 
13 

0414 
0792 
1139 

0453 
0828 
1173 

0492 
0864 
1206 

0531 
0899 
1239 

0569 
0934 
1271 

0607 
0969 
1303 

0645 
1004 
1335 

0682 
1038 
1367 

0719 
1072 
1399 

0755 
1106 
1430 

4   811 
3  710 
3  610 

15  19  23 
14  17  21 
13  16  19 

263034 
24  28  31 
23  26  29 

14 
15 
16 

1461 
1761 
2041 

1492 
1790 
2068 

1523 
1818 
2095 

1553 

1847 
2122 

1584 

1875 
2148 

1614 
1903 
2175 

1644 
1931 
2201 

1673 
1959 

2227 

1703 
1987 
2253 

1732 
2014 
2279 

369 
368 
358 

12  15  18 
11  14  17 
11  13  16 

21  24  27 
20  22  25 
18  21  24 

17 
18 
19 

20| 

2304 

2553 

2788 

2330 
2577 
2810 

2&55 
2G01 
2833 

2380 
2625 
2856 

2405 

2648 

2878 

2430 

2672 
2900 

2455 
2695 
2923 

2480 
2718 
2945 

2504 

2742 
2967 

2529 
2765 
2989 

257 
257 
247 

10  12  15 
9  1214 
9  11  13 

17  20  22 
161921 
101820 

3010 

3032 

3243 
3444 
3636 

3054 

3263 
3464 
3655 

3075 

3096 

3118 

3139 

3160 

3181 

3201 

246 

8  11  13 

15  17  19 

21 
22 
23 

3222 
3424 
3617 

3284 
3483 
3674 

3304 
3502 
3692 

3324 
3522 
3711 

3345 
3541 
3729 

3365 
3560 
3747 

3385 
3579 
3766 

3404 
3598 
3784 

246 
246 
246 

8  1012 
8  10  U 
7  9  11 

14  16  18 
141617 
13  15  17 

24 
25 
26 

3802 
3979 
4150 

3820 
3997 
4166 

3838 
4014 
4183 

3856 
4031 
4200 

3874 
4048 
4216 

3892 
4065 
4232 

3909 
4082 
4249 

3927 
4099 
4265 

3945 
4116 
4281 

3962 
4133 
4298 

245 
246 
235 

7  9  11 
7  9  10 
7  8  10 

12  14  16 
12  14  10 
11  13  15 

27 
28 
29 

4314 
4472 
4624 

4330 

4487 
4639 

4786 

4928 
5065 
5198 

4346 

4502 
4654 

4800 

4942 
5079 
5211 

4362 
4518 
4669 

4378 
4533 
4683 

4393 
4548 
4698 

4409 
4504 
4713 

4425 

4579 
4728 

4440 
4594 
4742 

4456 
4609 
4757 

235 
235 
134 

689 
689 
679 

11  12  14 
11  12  14 

10  12  13 

30 

4771 

4914 
5051 
5185 

4814 

4829 

4969 
5105 
5237 

4843 

4983 
5119 
5250 

4857 

4997 
5132 
5263 

4871 

4886 

4900 

134 

679 

10  11  13 

31 
32 
33 

4955 
5092 
5224 

5011 
5145 
5276 

5024 
5159 
5289 

5038 
5172 
6302 

134 
3   4 
3   4 

678 

578 
578 

10  11  12 
91112 
91112 

34 
35 

36 

5315 
5441 
5563 

5328 
5453 
5575 

5340 
5465 
5587 

5353 
5478 
5599 

5366 
5490 
5611 

5378 
5502 
5623 

5391 
5514 
5635 

5403 
5527 
5647 

6416 

5539 
5658 

5428 
5551 
6670 

2   4 
2   4 
2   4 

568 
567 
667 

91011 
91011 
81011 

37 
38 
39 

5682 
5798 
5911 

5694 
5809 
5922 

5705 
5821 
5933 

5717 
5832 
5944 

5729 
5843 
5955 

5740 
5855 
5966 

5752 

08C6 
5977 

5763 

5877 
50C8 

5775 

5888 
5999 

5786 
5899 
6010 

2   4 
2    3 
123 

567 
667 

457 

8   911 
8   910 
8    910 

40 

6021 

6031 

6042 

6053 

6064 

6075 

6085 

C096 

6107 

6117 

1   2    3 

456 

8    910 

41 

42 
43 

6128 
6232 
6335 

6138 
62  13 
6345 

6149 
6253 
6355 

6160 
6263 
6365 

6170 
6274 
6375 

6180 
6284 
C3S5 

6191 
6294 
6395 

6201 
6304 
6405 

6212 
6314 
6415 

6222 
6325 
6425 

123 
123 
123 

456 
466 
466 

789 
789 
789 

44 
45 
46 

6435 
6532 
6628 

6444 
6542 
6637 

6454 
6551 
6646 

6464 
6561 
6656 

6474 
6571 
6665 

6484 
6580 
6675 

6493 
6590 
6684 

6503 
6599 
6693 

6513 
6609 
6702 

6522 
6618 
6711: 

123 
123 
123 

456 
456 
456 

789 
789 

778 

47 
48 
49 

6721 
6812 
6902 

6730 
6821 
6911 

6739 
6830 
6920 

6749 
6839 
6928 

6758 
6848 
6937 

6767 
6857 
6946 

6776 
6866 
6955 

6785 
6875 
6964 

6794 

6884 
6972 

6803 
6893 
6981 

123 
123 
123 

456 
456 
445 

778 
778 
678 

j>P_ 

51 
52 
53 

6990 

7076 
7160 
7243 

6998 

7084 
7168 
7251 

7007 

7016 

7024 

7033 

7118 
7202 
7284 

7042 

7050 

7059 

7143 
7226 
7308 

7067 

7152 
7235 
7316 

123 

345 

678 

7093 
7177 
7259 

7101 

7185 
7267 

7110 
7193 
7275 

7126 
7210 
7292 

7135 

7218 
7300 

123 
123 
122 

346 
345 
345 

678 
677 
667 

54 

7324 

7332 

7340 

7348 

7356 

7364 

7372 

7380 

7388 

7396 

122 

345 

667 

N 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

122 

456 

789 

The  proportional  parts  are  stated  in  full  lor  every  tenth  at  the  right-hand  side. 
The  logarithm  of  any  number  of  four  significant  figures  can  be  read  directly  by  add- 


A] 


Table  A  —  Four  Place  Logarithms 


199 


IT 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

55 

56 

7404 
7482 

7412 
7490 

7419 
7497 

7427 
7505 

7435 
7513 

7443 
7520 

7451 

7528 

7459 
7536 

7466 
7543 

7474 
7551 

122 
122 

345 
345 

567 
567 

57 

58 
59 

7559 
7634 
7709 

7782 

7853 
7924 
7993 

7566 
7642 
7716 

7789 

7860 
7931 
8000 

7574 
7649 
7723 

7582 
7657 
7731 

7803 

7875 
7945 
8014 

7589 
7664 

7738 

7810 

7882 
7952 
8021 

7597 
7672 
7745 

7604 
7679 
7752 

7612 
7686 
7760 

7619 
7694 
7767 

7627 
7701 

7774 

112 
112 

112 

345 
344 
344 

567 
567 
567 

60 

61 
62 
63 

7796 

7818 

7825 

7832 

7839 

7846 

112 

344 

566 

7868 
7938 
8007 

7889 
7959 
8028 

7896 
7966 
8035 

7903 
7973 
8041 

7910 
7980 
8048 

7917 
7987 
8055 

112 
112 
112 

334 
334 
334 

566 
556 
556 

64 
65 

66 

8062 
8129 
8195 

8069 
8136 
8202 

8075 
8142 
8209 

8082 
8149 
8215 

8089 
8156 
8222 

8096 
8162 
8228 

8102 
8169 
8235 

8109 
8176 
8241 

8116 
8182 
8248 

8122 
8189 
8254 

112 
112 
112 

334 
334 
334 

556 
556 
556 

67 
68 
69 

8261 
8325 

8388 

8267 
8331 

8395 

8274 
8338 
8401 

8280 
8344 
8407 

8287 
8351 
8414 

8293 
8357 
8420 

8299 
8363 
8426 

8306 
8370 
8432 

8312 
8376 
8439 

8319 
8382 
8445 

112 
112 
112 

334 
334 
334 

556 
456 

456 

70 

8451 

8457 

8463 

8470 

8476 

8482 

8488 

8494 

8500 

8506 

112 

334 

456 

71 

72 
73 

8513 
8573 
8633 

8519 
8579 
8639 

8525 
8585 
8645 

8531 
8591 
8651 

8537 
8597 
8657 

8543 
8603 
8663 

8549 
8609 
8669 

8555 
8615 
8675 

8561 
8621 
8681 

8567 
8H27 
8686 

112 
112 
112 

334 
334 
2*4 

456 
456 
455 

74 
75 

76 

8692 
8751 
8808 

8698 
8756 
8814 

8704 
8762 
8820 

8710 

8768 
8825 

8716 
8774 
8831 

8722 
8779 
8837 

8727 
8785 
8842 

8733 
8791 
8848 

8739 
8797 
8854 

8745 
8802 
8859 

112 
112 
112 

234 
233 
233 

465 
455 
445 

77 
78 
79 

80 

8865 
8921 
8976 

8871 
8927 

8982 

8876 
8932 
8987 

8882 
8938 
8993 

8887 
8<)43 
8998 

8893 
8949 
9004 

8899 
8954 
9009 

8904 
89(50 
9015 

8910 

8965 
9020 

8915 
8971 
9025 

1  1  2 
112 
112 

233 
233 
233 

445 
445 
445 

9031 

9036 

9042 

9047 

9053 

9058 

9063 

9069 

9074 

9079 

1  1  2 

233 

445 

81 
82 
83 

9085 
9138 
9191 

9090 
9143 
9196 

9096 
9149 
9201 

9101 
9154 
9206 

9106 
9159 
9212 

9112 
9165 
9217 

9117 
9170 
9222 

9122 
9175 
9227 

9128 
9180 
9232 

9133 
9186 
9238 

1  1  2 
112 
112 

233 
233 
233 

445 
445 

445 

84 
85 

86 

9243 
9294 
9345 

9248 
92!)9 
9350 

9253 
9304 
9355 

9258 
9<09 
9360 

9263 
9315 
9365 

9269 
9320 
9370 

9274 
9325 
9375 

9279 
9330 
9380 

9284 
9335 
9385 

9289 
9340 
9390 

112 
112 
112 

233 
233 
233 

445 
445 
445 

87 
88 
89 

9395 
9445 
9494 

9400 
9450 
949!) 

9405 
9455 
9504 

9410 
9460 
9509 

9415 
9465 
9513 

9420 
9469 
9518 

9425 
9474 
9523 

9430 
9479 
9528 

9435 
9484 
9533 

9440 
9489 
9538 

112 
Oil 
0  1  1 

233 
223 
223 

445 
344 
344 

90 

9542 

9547 

9552 

9557 

9562 

9566 

9571 

9576 

9581 

9586 

Oil 

223 

344 

91 

92 
93 

9590 
9638 
9685 

9595 
9643 
9689 

9600 
9647 
9694 

9605 
9652 
9699 

9609 
9657 
9703 

9614 
9661 
9708 

9619 
9666 
9713 

9624 
9671 
9717 

9628 
9675 
9722 

9633 
9680 
9727 

Oil 
Oil 
Oil 

223 
223 
223 

3  4 
3  4 
3  4 

94 
95 

96 

9731 

9777 
9823 

9736 
9782 
9827 

9741 
9786 
9832 

9745 
9791 
9836 

9750 
9795 
9841 

9754 
9800 
9845 

9759 
9805 
9850 

9763 
9809 
9854 

9768 
9814 
9859 

9773 
9818 
9863 

Oil 
Oil 
Oil 

223 
223 
223 

3  4 
3  4 
3  4 

97 
98 
99 

9868 
9912 
9956 

9872 
9917 
9961 

9877 
9921 
9965 

9881 
9926 
9969 

9886 
9930 
9974 

9890 
9934 
9978 

9894 
9939 
9983 

9899 
9943 
9987 

9903 
9948 
9991 

9908 
9952 
9996 

Oil 
Oil 
Oil 

223 
223 
223 

3  4  4 
334 
334 

N 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

ing  the  proportional  part  corresponding  to  the  fourth  figure  to  the  tabular  numbei 
corresponding  to  the  first  three  figures.    There  may  be  an  error  of  1  in  the  last  place. 


200 


Table  B  —  Antilogarithms  to  Four  Places 


[B 


0 

1 

2 

8 

4 

5 

6 

7 

8 

9 

123 

456 

789 

.00 

1000 

1002 

1005 

1007 

1009 

1012 

1014 

1016 

1019 

1021 

001 

111 

222 

.01 
.02 
.03 

1023 
1047 
1072 

1026 
1050 
1074 

1028 
1052 
1076 

1030 
1054 
1079 

1033 
1057 
1081 

1035 
1059 
1084 

1038 
1062 
1086 

1040 
1064 
1089 

1042 
1067 
1091 

1045 
1069 
1094 

001 
001 
001 

111 
1  1 
1  1 

222 
222 
222 

.04 
.05 

.06 

.07 
.08 
.09 

.10 

.11 
.12 
.13 

1096 
1122 
1148 

1175 
1202 
1230 

1099 
1125 
1151 

1178 
1205 
1233 

1102 
1127 
1153 

1180 
1208 
1236 

1104 
1130 
1156 

1183 
1211 
1239 

1107 
1132 
1159 

1186 
1213 
1242 

1109 
1135 
1161 

1189 
1216 
1245 

1112 
1138 
1164 

1191 
1219 
1247 

1114 
1140 
1167 

1194 
1222 
1250 

1117 
1143 
1169 

1197 
1225 
1253 

1119 
1146 
1172 

1199 
1227 
1256 

Oil 
Oil 
Oil 

1  2 
1  2 
1  2 

222 
222 

222 

222 
223 

223 

Oil 
Oil 

112 
112 

1259 

1262 

1291 
1321 
1352 

1265 

1268 

1297 
1327 
135S 

1271 

1274 

1276 

1279 

1282 

1285 

Oil 

112 

223 

1288 
1318 
1349 

1294 
1324 
1355 

1300 
1330 
1361 

1303 
1334 
1365 

1306 
1337 
1368 

1309 
1340 
1371 

1312 
1343 
1374 

1315 
1346 
1377 

Oil 
Oil 
Oil 

122 
122 
122 

223 
223 
233 

.14 
.15 

.16 

1380 
1413 
1445 

1384 
1416 
1449 

1387 
1419 
1452 

1390 
1422 
1455 

1393 
1423 
1439 

1396 
1429 
1462 

1400 
1432 
1466 

1403 
1435 
1469 

1406 
1439 
1472 

1409 
1442 
1476 

Oil 
0  1  1 
Oil 

122 
122 
122 

233 
233 
233 

.17 

.18 
.19 

1479 
1514 
1549 

1483 
1517 
1552 

1486 
1521 
1556 

1489 
1524 
1560 

1493 

1528 
1563 

1496 
1531 
1567 

1500 
1535 
1570 

1503 
1538 
1574 

1507 
1542 
1578 

1510 
1545 
1581 

Oil 
Oil 
0  1  1 

122 
122 

1  2  2 

233 
233 
233 

.20 

.21 
.22 
.23 

1585 

1622 
1660 
1698 

1589 

1626 
1663 
1702 

1592 

1629 
1667 
1706 

1590 

1GOD 

1003 

1607 

1611 

1614 

1618 

0  1  1 

1  2  2 

333 

1633 
1071 
1710 

1637 
1675 
1714 

1641 
1679 
1718 

1644 
1683 
1722 

1648 
1687 
1726 

1052 
1690 
1730 

1656 
1694 
1734 

Oil 
Oil 
Oil 

122 
222 

222 

333 
333 
333 

.24 
.25 

.26 

1738 
1778 
1820 

1742 
1782 
1824 

1740 
1780 
1828 

1759 
1791 
1832 

1754 
1795 
1837 

1758 
1799 
1841 

1762 
1803 
1845 

1766 
1807 
1849 

1770 
1811 
1854 

1774 
1816 
1858 

Oil 
Oil 
Oil 

222 
223 

223 

334 
334 
334 

.27 

.28 
.29 

1862 
1905 
1950 

1866 
1910 
1954 

1871 
1914 
1959 

1875 
1919 
1963 

1879 
1923 
1968 

1881 
1928 
1972 

1888 
1932 
1977 

1892 
1936 
1982 

1897 
1941 

1986 

1901 
1945 
1991 

Oil 
Oil 
Oil 

223 
223 
223 

334 
344 
344 

.30 

.31 
.32 
.33 

1995 

2000 

2004 

2009 

2014 

2018 

2023 

2028 

2032 

2037 

Oil 

223 

344 

2042 

2089 
2138 

2046 
2094 
2143 

2051 
2099 
2148 

2056 
2104 
2153 

2061 
2109 
2158 

2065 
2U3 
2163 

2070 
2118 
2168 

2075 
2123 
2173 

2080 
2128 
2176 

2084 
2133 
2183 

Oil 
Oil 
Oil 

223 
223 
223 

344 
344 
344 

.34 
.35 

.36 

2188 
223!) 
2291 

2193 
2244 
2296 

2198 
22i9 
2301 

2203 
2254 
2307 

2208 
2259 
2312 

2213 
2235 
2317 

2218 
2270 
2323 

2223 
2275 
2328 

2228 
2280 
2333 

2234 
2286 
2339 

112 
112 

1  1  2 

233 
233 
233 

445 

445 
445 

.37 
.38 
.39 

.40 

2344 
2399 
2455 

2350 
2404 
24(50 

2355 
2410 
2466 

2360 
2415 

2472 

2366 
2421 
2477 

2371 

2427 
2483 

2377 
2432 
2489 

2382 
2438 
2495 

2388 
2443 
2500 

2393 
2449 
2506 

112 
112 
112 

233 
233 
233 

445 
455 
455 

2')12 

2518 

2523 

2529 

2535 

2594 
2655 
2716 

2541 

2547 

2553 

2559 

2564 

112 

234 

455 

.41 
.42 
.43 

2570 
2630 
2692 

2576 
2636 
2698 

2582 
2612 
2704 

2588 
2649 
2710 

2600 
2661 
2723 

2606 
2667 
2729 

2612 
2673 
2735 

2618 
2679 
2742 

2624 
2685 
2748 

112 
112 
112 

234 
234 
234 

456 
456 
456 

.44 
.45 
.46 

2754 
2818 
2884 

2761 
2825 
2891 

2767 
2831 
2897 

2773 
2838 
2904 

2780 
2844 
2911 

2786 
2851 
2917 

2793 
2858 
2924 

2799 
2864 
2931 

2805 
2871 
2938 

2812 
2877 
2944 

112 
112 
112 

334 
334 
334 

456 
556 
556 

.47 
.48 
.49 

2951 
3020 
3090 

2958 
3027 
3097 

2965 
3034 
3105 

2972 
3041 
3112 

2979 
3048 
3119 

2985 
3055 
3126 

2992 
3062 
3133 

2999 
3069 
3141 

3006 
3076 
3148 

3013 
3083 
3155 

112 
112 
112 

334 
334 
344 

566 
566 
566 

B] 


Table  B — Antilogarithms  to  Four  Places 


20 1 


0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

123 

456 

789 

.50 

3162 

3170 

3177 

3184 

3192 

3199 

3206 

3214 

3221 

3228 

567 

.51 
.52 
.53 

3236 
3311 
3388 

3243 
3319 
3396 

3251 
3327 
3404 

3258 
33M 
3412 

3266 
3342 
3420 

3273 
3350 
3428 

3281 
3357 
3436 

3289 
3365 
3443 

3296 
3373 
3451 

3304 
3381 
3459 

112 
1  1  2 
122 

344 
345 
345 

567 
567 
667 

.54 
.55 

.56 

3467 
3548 
3631 

3475 
3556 
3639 

3483 
3565 
3648 

3491 
3573 
3656 

3499 
3581 
3664 

3508 
3589 
3673 

3516 
3597 
3681 

3524 
3606 
3690 

#532 
3614 
3698 

3540 
3622 
3707 

122 
122 
122 

345 
345 
345 

667 
677 
678 

.57 
.58 
.59 

3715 
3802 
3890 

3724 
3811 

3899 

3733 
3819 
3908 

3741 
3828 
3917 

3750 
3837 
3926 

3758 
3846 
3936 

3767 
3855 
3945 

3776 
3364 
3954 

3784 
3873 
3963 

3793 

3882 
3972 

123 
123 
123 

345 
345 
455 

678 
678 
678 

.60 

.61 
.62 
.63 

3981 

3990 

3999 

4009 

4018 

4027 

4036 

4046 

4055 

4064 

123 

456 

788 

4074 
4169 
426(i 

4083 
4178 
4276 

4093 
4188 

4285 

4102 
4198 
4295 

4111 
4207 
4305 

4121 
4217 
4315 

4130 
4227 
4325 

4140 
4236 
4335 

4150 
4246 
4345 

4159 
42f>6 
4355 

123 
123 
123 

456 
456 
456 

789 
789 
789 

.64 
.65 

.66 

43fi5 
4467 
4571 

4375 
4477 
4581 

4385 
4487 
4592 

4395 

4498 
4603 

4406 
4508 
4613 

4416 
4519 
4624 

4426 
452!) 
4634 

4436 
4539 
4645 

4446 
4550 
4656 

4457 
4560 
4667 

123 
123 
123 

456 
456 
456 

789 
789 
7   910 

.67 
.68 
.69 

.70 

4677 
4786 
4898 

4688 
4797 
4909 

5023 

4699 

4808 
4920 

4710 
4819 
4932 

4721 
4831 
4943 

4732 
4842 
4955 

4742 
4853 
4966 

4753 
4864 
4977 

4764 
4875 
4989 

4775 

4887 
5000 

123 
123 
123 

457 
567 
567 

8   910 

8   910 
8    910 

5012 

5035 

5047 

5058 

5070 

5082 

5093 

5105 

5117 

123 

567 

8   910 

.71 
.72 
.73 

5129 
5248 
5370 

5140 
5260 
5383 

5152 
5272 
5395 

5164 
5284 
5408 

5176 

5297 
5420 

5188 
5309 
5433 

5200 
5321 
5445 

5212 
5333 
5458 

5224 
5346 
5470 

5236 
5358 
5483 

124 
1  2  4 
134 

567 
567 
567 

81011 
91011 
91011 

.74 
.75 

.76 

5495 
5623 
5754 

5508 
5636 
5768 

5521 
5649 
5781 

5534 
5662 
5794 

5546 
5675 
5808 

5559 
5689 
5821 

5572 
5702 
5834 

5585 
5715 
5848 

5598 
5728 
5861 

5610 
5741 
5875 

134 
134 
134 

568 
578 
578 

91012 
911  12 
91112 

.77 
.78 
.79 

5888 
6026 
6166 

5902 
6039 
6180 

5916 
6053 
6194 

5929 
6067 
6209 

5943 
6081 
6223 

5957 
60!  )5 
6237 

5970 
6109 
6252 

5984 
6124 
626(5 

5998 
6138 
6281 

6012 
6152 
6295 

134 
134 
134 

578 
678 
679 

10  11  12 
10  11  13 
10  11  13 

.80 

6310 

6324 

6339 

6353 

6368 

6383 

6397 

6412 

6427 

6442 

134 

679 

10  12  13 

.81 
.82 
.83 

6457 
6607 
6761 

6471 
6622 
6776 

6486 
6637 
6792 

6501 
6653 
6808 

6516 
6668 
6823 

6531 
6683 
6839 

6546 
6699 
6855 

6561 
6714 
6871 

6577 
6730 
6887 

6592 
6745 
6902 

235 
235 
235 

689 
689 
689 

11  1214 
11  12  14 
11  13  14 

.84 
.85 

.86 

6918 
7079 
7244 

6934 
7096 
7261 

6950 
7112 

7278 

6966 
7129 
7295 

6982 
7145 
7311 

6998 
7161 
7328 

7015 
7178 
7345 

7031 
7194 
7362 

7047 
7211 
7379 

7063 

7228 
7396 

235 
235 
235 

7  810 
7  810 
7  810 

11  13  15 
12  13  15 
12  14  15 

.87 
.88 
.89 

.90 

.91 
.92 
.93 

7413 

7586 
7762 

7943 

8128 
8318 
8511 

7430 
7603 
7780 

7962 

8147 
8337 
8531 

7447 
7621 
7798 

7980 

8166 
8356 
8551 

7464 
7(538 
7816 

7482 
7656 
7834 

8017 

7499 
7674 
7852 

8035 

7516 
7691 
7870 

80.54 

7534 
7709 
7889 

8072 

7551 

7727 
7907 

8091 

7568 
7745 
7925 

8110 

245 
245 
246 

7  910 

7  911 
7  911 

12  14  16 
12  14  16 
13  15  1(5 

7998 

8185 
8375 
8570 

246 

7  911 

13  15  17 

8204 
8395 
8590 

8222 
8414 
8610 

8241 
8433 
8630 

8260 
8453 
8650 

8279 
8472 
8670 

8299 
8492 
8690 

246 
246 
246 

8  911 
81012 
81012 

13  15  17 
14  15  17 
14  16  18 

.94 
.95 

.96 

8710 
8913 
9120 

8730 
8933 
9141 

8750 
8954 
9162 

8770 
8974 
9183 

8790 
8995 
9204 

8810 
9010 
9226 

8831 
90136 
9247 

8851 
9057 
9268 

8872 
9078 
9290 

8892 
9099 
9311 

246 
246 
246 

81012 
8  10  12 
911  13 

14  16  18 
15  17  19 
15  17  19 

.97 

.98 
.99 

9333 
9550 
9772 

9354 
9572 
9795 

9376 
9594 
9817 

9397 
9616 
9840 

9419 
96.38 
9863 

9441 
9661 
9886 

9462 
9683 
9908 

9484 
9705 
9931 

9506 
9727 
9954 

9528 
9750 
9977 

246 
247 
257 

911  13 
91113 
91114 

15  17  19 
16  18  20 
16  18  21 

INDEX 


Acidimetry 105,  115 

Acids,  degree  of  ionization  of  .     .  24 

determination  of  the  neutraliza- 
tion value  of 115 

standard  solutions  of  .     .     .     .  105 

titration  of 108 

Accuracy 3>  4 

Adsorption 22,  30 

Affinity  constant 25 

Afterflow,  error  from 45 

Alkalimeter,  Mohr's 78 

Alkalimetry 105,  108,  113 

Alkali  solutions,  standard    ...  105 

Aluminum,  determination  of    .     .  62 
Ammonium  thiocyanate,  standard 

solutions  of 156 

Ampere,  definition  of 90 

Analyzed  chemicals 6 

Antilogarithms 200 

Antimony,    determination    of    in 

stibnite 148 

Apparatus,  list  of  for  quantitative 

work 197 

Arsenic,  removal  of  from  copper  153 

Arsenious  oxide,  primary  standard  147 

Asbestos,  preparation  of  for  niters  33 

use  of  in  nitration  .  . 
Ashless  filter  papers  .  . 
Atomic  weights,  table  of 


-     33,  34 
.     -       30 
Back  cover 
sheet 


Baking  powder,  determination  of 

carbon  dioxide  in      ....  79 

Balance,  analytical 7 

adjustment  of 9 

conditions  to  be  fulfilled  by  .     .  8 

exercises  with S3 

location  of 9 

relative  length  of  arms  of     .     .  17 

sensitivity  of 13 

use  and  care  of 9 

zero-point  of 1 1 

Barium  sulphate,  properties  of      63,  64 


Bases,  degree  cf  ionization  of  .     .  24 

standard  solutions  of  .     .     .     .  105 

titration  of 108 

Bumping 42 

Buoyancy,  correction  for     ...  18 

Burettes 44 

calibration  of 49 

cleaning  of 46 

reading  of 46 

Burning  filter  papers 38 

Calcium,  determination  cf  in  lime- 
stone       70,  133 

oxalate,  properties  cf  .     .     .     73,  74 
Calibration  of  measuring  vessels  .      49 

of  weights 14 

Carbon  dioxide,  determination  of 

in  limestone 76 

determination    of    in    baking 

powders 79 

Chemical  equilibrium  ...  23,  24 
Chemical  factors  ....  161,  162 
Chlorine,  determination  of  54,  59,  158 
Chrome  iron  ore,  determination  of 

chromium  in 125 

Chromic  oxide,  ignition  of  ...  63 
Chromium,  determination  of  .  62,  125 
Cleaning  solution,  preparation  and 

use  of 196 

Colloidal  precipitates .     ....       21 

Common-ion  effect 26 

Contamination  of  precipitates .     .       22 

Contat-Gockel  valve 135 

Copper,  determination  of  .  .  86,  152 
Counterpoise,  use  of  in  weighing 

bulky  objects 18,  80 

Crucibles,  materials  of    ....      40 

Current  density 94 

Current,  production  of  for  electro- 
analysis      90 

Decantation 32 

Decimal,    number    of    places    to 

report 5 


203 


204 


INDEX 


Deposition  voltages  of  elements, 

table  of 92 

Desiccators ,  ;'.    .  39 

Dichromate  processes      ....  120 

solutions,  standard 121 

Dichromate-sulphuric  acid  clean- 
ing solution 196 

Digestion  of  precipitates      ...  21 

Distilled  water,  testing  of    ...  6 

Double  precipitation 23 

Drainage,  error  from 45 

Drying  ovens 37 

Economy  of  time 4 

Electro-analysis 86 

Electrode  potentials,  table  of  .     .  92 
Electrodes,  material  and  form  of  .  95 
Electrolytes,  influence  of  composi- 
tion of  upon  electro-analysis .  94 
Electrolytic  separations  ....  92 
Electrolytic  solution  tension    .     .  91 
End-point  in  titration,  determina- 
tion of loo 

Equilibrium,  chemical     .     .     .     23,  24 

Equilibrium  constant      ....  25 

Evaporation  of  liquids    ....  41 

Factor,  chemical    ....      161,  162 

normality 164 

Faraday's  laws  .     .......  90 

Ferric  alum,  indicator     .     .      156,  157 

Ferric  oxide,  ignition  of  ....  62 

Ferrous      ammonium      sulphate, 

standard  solutions  of    .     .     .  121 
Fertilizers,  determination  of  nitro- 
gen in    116 

Filters,  selection  and  use  of      .    .  30 

Filtrates,  testing  of 32 

Filtration 30 

Fine-grained  precipitates,  enlarge- 
ment of  the  particles  of    .     .  21 
Flasks,  volumetric      .....  44 

calibration  of 49 

Flocculation  of  colloids  ....  22 

Funnels,  selection  of 31 

Fusions,  removal  of  from  crucibles  81 

Gelatinous  precipitates   ....  32 
Gooch  filters,  preparation  and  use 

of 33,  34 

sources  of  error  with    ....  35 


Graduated  cylinders 45 

Gravimetric  analysis i,  53 

Guyard's  reaction 139 

Halogens,  determination  of  ...      59 
Hematite,  determination  of  iron 

in 130 

Hydrochloric  acid,  standard  solu- 
tions of 109 

Ignition  of  precipitates   ....      37 

Indicators 100,  107,  120 

general  theory  of 101 

sensitiveness  of  in  alkalimetry 

and  acidimetry  ....  105, 107 
Indirect  methods  of  analysis  .  .  164 
Insoluble  matter,  determination  of 

in  limestone 71,  72 

International      atomic     weights, 

table  of  ...  Back  cover  sheet 
Iodine,  standard  solutions  of  .  146, 147 

lodometric  processes 142 

lonization,  degree  of 24 

repression  of 24 

Ions,  complex 24 

composition  of  the 24 

Iron,  determination  of  ....  59 
oxidation  of  ferrous  to  ferric  61,  122 
reduction  of  ferric  to  ferrous 

62,  123,  132, 137 

Iron  ores,  decomposition  of  124, 131, 132 
Iron  wire,  primary  standard     .     .     123 

Jones  reductor,  assembly  and  use 

of 137 

Kjeldahl,   determination  of  pro- 
tein nitrogen 116 

Labels 4 

Lead,  determination  of  in  an  ore  150 

Limestone,  analysis  of  .  .  70,  76,  133 
determination  of  carbon  dioxide 

in 76 

Liquids,  evaporation  of  ....  41 

transference  of    .     .    .    fc    .    .  42 

volumetric  measurement  of  .     .  43 

Liter,  Mohr's     .......  48 

normal 48 

true 48 


INDEX 


205 


Limits  of  error  in  experimental 

work 4 

Logarithms 198 

Magnesium,  determination  of  in 

limestone 70,  75 

Magnesium     ammonium     phos- 
phate, ignition  of     ...     68,  70 
Manganese,  determination  of  in 

an  ore .  139 

Manganese  ores,  decomposition  of  141 

Methyl  orange  solution  ....  109 

Neatness 3 

Neutral  solution,  definition  of  .  .  106 
Neutralization  methods  .  .  .  105, 108 
Nickel,  electrolytic  determination 

of 89 

Nitrogen,  Kjeldahl  determination 

of .     . 116 

Normality  factor,  definition  of      .164 

Normal  solutions 98 

Normal  System  of  Reagents,  ad- 
vantages of 193 

Notebooks 4 

Ohm,  definition  of 90 

Ohm's  law 90 

Ovens,  drying 37 

Overvoltage       92 

Oxidation  and  reduction  methods  119 
Oxidizing  agents,  standard  solu- 
tions of 119 

Parallax,  error  from 46 

Permanganate  processes .     .     .     .  127 
Permanganate  solutions,  standard  128 
Phenolphthalein  solution     .     .     .  109 
Phosphoric  anhydride,  determina- 
tion of 66 

Phosphorus,  determination  of  in 

steel 135 

Pipettes,  transfer 44 

calibration  of 49 

Platinum  ware,  defects  of  mod- 
ern      41 

specifications  for 41 

use  and  care  of 41,  87 

Polarity  of  terminals,  determina- 
tion of 87 

Polarization 92 

Policeman,  definition  of  ....  32 

Potash,  determination  of     ...  84 


Potassium     bitartrate,     primary 

standard 113 

dichromate,  standard  solutions 

of 121 

ferricyanide,  indicator      .     .120,121 
iodate,  primary  standard        143,  148 
permanganate,    standard    solu- 
tions of 128 

thiocyanate,  standard  solutions 
of       .156 

Precipitates,  colloidal      ....  21 

contamination  of 22 

digestion  of 21 

drying  of 37 

enlargement  of  the  particles  of .  21 

fine-grained 21 

flocculation  of 22 

for  use  in  gravimetric  analysis  20 

ignition  of 37 

purification  of 23 

washing  of 22, 30 

Precipitation 20 

theory  of 23 

volumetric  methods  of     ...  155 

Problems 166 

Pyrolusite,  oxidizing  value  of  .     .  134 

Questions 178 

Reaction,  Guyard's 139 

Reactions  suitable  for  use  in  volu- 
metric analysis 99 

Reagents,  analyzed 6 

preparation  of 193 

quality  of 6 

testing  of 6 

Records 4,  5 

Reducing  agents,  standard  solu- 
tions of 119 

Reduction,  methods  of  oxidation 

and 119 

Reductor,  Jones     .     .    .    .    .    .  137 

Reversible  reactions   .....  23 

Salts,  degree  of  ionization  of    .     .  24 
Samples,  Preparation  of,  for  An- 
alysis      51,  196 

Saturated  solution,  definition  of  .  28 

Sensitivity,  of  balance     ....  13 

of  indicators,  table  of  the     .     .  107 

Siderite,  determination  of  iron  in  124 

Silica,  determination  of  ....  80 


2O6 


INDEX 


Silicates,  determination  of  silica  in 

refractory 80 

Silicic  acid,  dehydration  of  ...  83 
Silver,  determination  of  .  59,  155,  158 
Silver,  primary  standard  .  .  .  157 
Silver  chloride,  properties  of  .  57,  58 
Silver  ion,  properties  of  ....  58 
Silver  nitrate,  standard  solutions 

of 156 

Soda  ash,  alkaline  value  of  .     .     .     113 
Sodium  carbonate,  primary  stan- 
dard   109,  113 

Sodium  chloride,  determination  of 

chlorine  in 54 

primary  standard 157 

purification  of 157 

Sodium  hydroxide,  standard  solu- 
tions of 109 

Sodium  oxalate,  primary  standard  129 
Sodium  thiosulphate,  standard 

solutions  of     ...  146,  151,  152 
Solubility,  effect  of  size  of  par- 
ticles on 21 

Solubility  product 27 

Solution  tension 27 

electrolytic 91 

Solution  of  iron  ores  .     .   124,  131,  132 

Solution  of  manganese  ores      .     .     141 

Standardization,  definition  of  .     .      97 

of  hydrochloric  acid     .     .     .     .     109 

of  sodium  hydroxide  solution    .     109 

of  dichromate  solution     .     .     .     121 

of  ferrous  ammonium  sulphate 

solution 121 

of  permanganate  solution     .     .     128 

of  iodine  solution 146 

of    sodium    thiosulphate    solu- 
tion   146,  151,  152 

of  silver  nitrate  solution  .     .     .     156 
of  thiocyanate  solution    .     .     .     156 
Standard  solution,  definition  of    .      97 
Search,  indicator    .     .     .     .      144,  145 
Stibnite,  determination  of   anti- 
mony in 148 

determination  of  sulphur  in  .     .     163 

Stoichiometry 159 

Suction,  Mse  of 31,  33,  34 

Sulphur,  determination  of   .    .     63,  65 


Temperature,  correction  for  dif- 
ferences in 46 

Tension,  solution    .     .     ....     .  27 

electrolytic  solution     ....  91 

Testing  for  complete  precipitation  32 

of  washings 32 

Thiocyanate  solutions,  standard  .  156 

Titration,  definition  of    ....  2 

Transference  of  liquids    ....  42 

Transfer  pipettes 44 

Triangles .     38, 41 

Vacuum,  use  of      ....    31,  33,  34 

Valve,  Contat-Gockel     .     .     .     .  135 

Volt,  definition  of 90 

Volume,  units  of    ......  48 

Volumetric  analysis,  general  dis- 
cussion        97 

neutralization  methods  of     .     .  105 
oxidation        and        reduction 

methods  of 119 

precipitation  methods  of  ...  155 

reactions  suitable  for  ....  99 

Volumetric  apparatus      ....  44 

calibration  of 49 

cleaning  of 46 

necessary  precautions  in  the  use 

of 45 

Volumetric   System,   Advantages 

of     ^ ^.   43,  103 

Volumetric  Work,  General  Direc- 
tions       ....  103 

Wash  bottles 33 

Washing  of  precipitates  .     .    22,  30,  32 

theory  of 35 

Washings,  testing  of  .    .    .'   .    .  32 

Water,  ionization  of   .     .     .     .     .  106 

Weighing ,    .  7,  53 

limits  of  error  in 1 1 

methods  of 12 

summary  of 20 

Weights,  calibration  of    ....  14 

use  and  care  of 10 

Zero-point  of  balance,  determina- 
tion of ii 

Zimmermann-Reinhardt  solution .  131 


Printed  in  the  United  States  of  America. 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 


Return  to  desk  from  which  borrowed. 
This  book  is  DUE  on  the  last  date  stamped  below. 

ENGINEERING  LIBRA 


MAR  24 


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YC  o 


UNIVERSITY  OF  CALIFORNIA 

DEPARTMENT    OF    CIVIL    ENGINEERING 

f  PRi    SLEY.  C*     IFORNIA 

INTERNATIONAL  ATOMIC  WEIGHTS,  1917 


Aluminum 

Al 

27.1 

Molybdenum 

Mo 

96.0 

Antimony 

Sb 

1  20.  2 

Neodymium 

Nd 

144-3 

Argon 

A 

39-88 

Neon 

Ne 

20.2 

Arsenic 

As 

74.96 

Nickel 

Ni 

58.68 

Barium 

Ba 

137-37 

Niton 

Nt 

222-4 

Bismuth 

Bi 

208.0 

Nitrogen 

N 

I4.OI 

Boron 

B 

II.O 

Osmium 

Os 

I9I.9 

Bromine 

Br 

79.92 

Oxygen 

0 

I6.OOO 

Cadmium 

Cd 

112.40 

Palladium 

Pd 

106.7 

Caesium 

Cs 

132.81 

Phosphorus 

P 

31.04 

Calcium 

Ca 

40.07 

Platinum 

Pt 

195.2 

Carbon 

C 

12.005 

Potassium 

K 

39.10 

Cerium 

Ce 

140.25 

Praseodymium 

Pr 

140.9 

Chlorine 

Cl 

35.46 

Radium 

Ra 

226.0 

Chromium 

Cr 

52.0 

Rhodium 

Rh 

IO2-9 

Cobalt 

Co 

58.97 

R.  oidium 

Rb 

8545 

Columbium 

Cb 

93-i 

Rii  '  henium 

Ru 

IOI-7 

Copper 

Cu 

63-57 

S;.  iiiarium 

Sa 

150.4 

Dysprosium 

Dy 

162.5 

Scandium 

Sc 

44.1 

Erbium 

Er 

167.7 

Selenium 

.    Se 

79.2 

Europium 

Eu 

152.0 

Silicon 

Si 

28.3 

Fluorine 

F 

19.0 

Silver 

Ag 

IO/.88 

Gadolinium 

Gd 

157.3 

Sodium 

Na 

23.00 

Gallium 

Ga 

69.9 

Strontium 

Sr 

87.63 

Germanium 

Ge 

72-5 

Sulphur 

S 

32.06 

Glucinum 

Gl 

9.1 

Tantalum 

Ta 

l8l.5 

Gold 

Au 

197.2 

Tellurium 

Te 

127-5 

Helium 

He 

4.00 

Terbium 

Tb 

159.2 

Holmium 

Ho 

163-5 

Thallium 

Tl 

204.0 

Hydrogen 

H 

1.008 

Thorium 

Th 

232.4 

Indium 

In 

114.8 

Thulium 

Tm 

168.5 

Iodine 

I 

126.92 

Tin 

Sn 

II8.7 

Iridium 

Ir 

I93-I 

Titanium 

Ti 

48.1 

Iron 

Fe 

55.84 

Tungsten 

W    • 

184.0 

Krypton 

Kr 

82.92 

Uranium 

U 

238.2 

Lanthanum 

La 

139.0 

Vanadium 

V 

51.0 

Lead 

Pb 

207.20 

Xenon 

Xe 

130.2 

Lithium 

Li 

6.94 

Ytterbium 

Yb 

173-5 

Lutetium 

Lu 

175-0 

Yttrium 

Yt 

88.7 

Magnesium 

Mg 

24.32 

Zinc 

Zn 

65.37 

Manganese 

Mn 

54-93 

Zirconium 

Zr 

9O.6 

Mercury 

Hg 

200.6 

